17Electrochemistry

📚
ELECTROCHEMISTRY
▢ Definition: Branch of chemistry studying relationship between electrical energy and chemical change.
▢ Electrical Cells: Convert electrical energy into chemical energy.
▢ Coulometer: Apparatus used for measuring quantity of electricity.
▢ Conductors:

Table 1: Types of conductors

Type
Meaning
Charge carrier
Examples
Temperature effect
Metallic / electronic conductor
Current flow without chemical decomposition
Electrons
Cu, Ag, Al, Pt, graphite, alloys
\(Conductance \propto \dfrac{1}{Temperature}\)
Electrolytic conductor
Aqueous / molten electrolyte allows current with chemical decomposition
Ions
Acids, bases, salts in solution or fused state
\(Conductance \propto Temperature\)
Non-electrolyte
Solution / molten state does not conduct electricity
No free ions
Organic compounds, pure water
📚
METALLIC VS ELECTROLYTIC CONDUCTION

Table 1: Distinction between metallic and electrolytic conduction

Feature
Metallic / electronic conduction
Electrolytic conduction
Substance decomposition
No decomposition
Decomposition occurs
Current carrier
Electrons
Ions
Matter flow
No flow of matter
Matter transfer as ions
Chemical change
Absent
Oxidation/reduction at electrodes
Temperature effect on resistance
Resistance increases with temperature
Resistance decreases with temperature
Faraday's law
Not followed
Followed
Examples
Metals, graphite, minerals
\(H_2SO_4\), \(NaCl\), \(KOH\), \(CH_3COOH\) solutions
Important
  • •
    Ohm's law is followed by metallic and electrolytic conductors.
  • •
    Faraday's law is followed only by electrolytic conductors.
  • •
    Anode = oxidation electrode.
  • •
    Cathode = reduction electrode. [MOE]
📚
ELECTROLYSIS
▢ Definition: Decomposition of electrolyte by passing electric current through aqueous solution or fused state.
▢ Ion Movement:
  • •
    Cations move to cathode → gain electrons → neutral atoms.
  • •
    Anions move to anode → lose electrons → neutral atoms.
  • •
    If more than one ion attracted to same electrode, ion requiring least energy / having highest discharge potential is discharged first.
▢ Discharge Series:

Table 1: Decreasing discharge potential / increasing decomposition order

Ion type
Series
Cations
\(Li^+, K^+, Ca^{2+}, Na^+, Mg^{2+}, Al^{3+}, Zn^{2+}, Fe^{2+}, Ni^{2+}, H^+, Cu^{2+}, Hg^{2+}, Ag^+, Au^{3+}\)
Anions
\(SO_4^{2-}, NO_3^-, OH^-, Cl^-, Br^-, I^-\)
▢ Aqueous Solution Rules:
  • •
    Cation left of \(H^+\) → \(H^+\) from water reduced at cathode, not that cation.
  • •
    Anion left of \(OH^-\), e.g. \(SO_4^{2-}\), \(NO_3^-\) → \(OH^-\) from water oxidized at anode.
  • •
    Cation with higher reduction potential deposits at cathode first.
  • •
    Anion with low reduction potential / high oxidation potential oxidizes at anode first.
▢ Device: Electrolytic cell / voltameter; not voltmeter.
▢ Electrolysis of Water: During electrolysis of water with dilute \(H_2SO_4\), concentration of \(H_2SO_4\) increases with time due to decrease in water volume.
📚
NATURE OF RESULTING SOLUTION AFTER ELECTROLYSIS

Table 1: Resulting solution nature

Condition
Result
Examples
Cation and anion left of \(H^+\) and \(OH^-\)
Neutral
\(Na_2SO_4\), \(KNO_3\)
Cation and anion right of \(H^+\) and \(OH^-\)
Neutral
\(CuCl_2\), \(AuCl_3\)
Anion right of \(OH^-\), cation left of \(H^+\)
Basic
\(NaCl\), \(KCl\), \(CaCl_2\)
Anion left of \(OH^-\), cation right of \(H^+\)
Acidic
\(CuSO_4\), \(AgNO_3\)
📚
ELECTROPLATING AND ELECTROREFINING
▢ Electroplating:
  • •
    Article to be plated = cathode.
  • •
    Precious metal = anode.
  • •
    Electrolyte contains cation of precious metal.
  • •
    Copper electroplating: article as cathode, copper as anode, \(CuSO_4\) as electrolyte.
  • •
    If electrode is active at cathode, metal deposits on cathode and anode metal dissolves.
▢ Electrorefining:
  • •
    Pure metal = cathode.
  • •
    Impure sample = anode.
  • •
    Example: impure copper purification by electrolysis.
📚
FARADAY'S LAWS OF ELECTROLYSIS
▢ First Law:
❖ Statement: Mass of ion discharged during electrolysis is directly proportional to quantity of electricity passed.
❖ Formulae:
  • •
    \(W \propto Q\)
  • •
    \(W=ZQ\)
  • •
    \(Q=It\)
  • •
    \(W=ZIt\)
❖ Terms:
  • •
    \(W=\) mass deposited/liberated
  • •
    \(Z=\) electrochemical equivalent
  • •
    \(I=\) current in ampere
  • •
    \(t=\) time in second
❖ Electrochemical Equivalent: Mass of ion deposited by passing current of 1 ampere for 1 second.
❖ Unit of \(Z\): Gram per coulomb.
▢ Second Law:
❖ Statement: Same quantity of electricity passed through different electrolytes gives masses liberated at electrodes directly proportional to their chemical equivalents.
❖ Formulae:
  • •
    \(W \propto E\)
  • •
    \(\dfrac{W}{E}=constant\)
  • •
    \(\dfrac{W_1}{W_2}=\dfrac{E_1}{E_2}\)
  • •
    \(\dfrac{Z_1}{Z_2}=\dfrac{E_1}{E_2}\)
  • •
    \(Z \propto E\)
  • •
    \(Z=\dfrac{E}{F}\)
❖ Terms:
  • •
    \(E=\) equivalent weight / chemical equivalent
  • •
    \(F=\) Faraday constant
▢ Faraday:
  • •
    \(1\,Faraday = 96500\,C\) [MOE 2064]
  • •
    Unit of Faraday = coulomb/mole.
  • •
    Charge carried by 1 mole electrons = 1 Faraday.
  • •
    Quantity of charge liberating 1 gram equivalent of each element = 1 Faraday.
  • •
    \(1F = 96500\,C = 6.023\times10^{23}\) electrons.
  • •
    \(\text{Number of equivalents}=\dfrac{\text{mass deposited/liberated}}{\text{equivalent wt.}}=\dfrac{\text{coulombs}}{96500}=\text{number of Faradays}\)
▢ One Faraday Deposits:

Table 1: Moles deposited by 1F

Cation valency
Moles deposited by 1F
Monovalent
1 mole
Divalent
\(\dfrac{1}{2}\) mole
Trivalent
\(\dfrac{1}{3}\) mole
n-valent
\(\dfrac{1}{n}\) mole
📚
QUALITATIVE ASPECTS OF ELECTROLYSIS
▢ Products Depend On:
  • •
    Nature of electrolyte
  • •
    Nature of electrode
  • •
    State and concentration of electrolyte
▢ Nature of Electrolyte:

Table 1: Products with inert electrode

Electrolyte
Cathode product
Anode product
Electrode
\(CuSO_4(aq)\)
Cu
\(O_2\)
Inert: graphite / platinum / gold
\(AgNO_3(aq)\)
Ag
\(O_2\)
\(NaCl(aq)\)
\(H_2\)
\(Cl_2\)
KI
\(H_2\)
\(I_2\)
\(KNO_3\)
\(H_2\)
\(O_2\)
HBr
\(H_2\)
\(Br_2\)
NaOH
\(H_2\)
\(O_2\)
❖ Rule:
  • •
    Greater reduction potential → cathode product.
  • •
    Greater oxidation potential → anode product.
▢ Nature of Electrode:

Table 1: Effect of electrode nature

Electrolyte
Electrode
Products / change
\(AgNO_3(aq)\)
Inert electrode
Cathode: Ag; anode: \(O_2\)
\(AgNO_3(aq)\)
Ag electrode
Cathode: Ag; anode: \(Ag^+\)
\(CuSO_4(aq)\)
Inert electrode
Cathode: Cu; anode: \(O_2\)
\(CuSO_4(aq)\)
Cu electrode
Cathode: Cu; anode: \(Cu^{2+}\)
\(CuSO_4(aq)\)
Ag electrode
Cu deposits on Ag
\(AgNO_3(aq)\)
Cu electrode
Ag deposits on Cu
Metal oxidation
Oxidation of metals lying below hydrogen is generally more favourable than oxidation of water to \(O_2\).
📚
CONDUCTANCE
▢ Conductance:
  • •
    \(C=\dfrac{1}{R}\)
  • •
    \(R=\) resistance
  • •
    Unit of \(R\): ohm \((\Omega)\)
  • •
    Unit of conductance: \(ohm^{-1}\), mho, Siemens
▢ Specific Conductance / Conductivity:
❖ Relations:
  • •
    \(E=IR\)
  • •
    \(R=\rho\dfrac{l}{a}\)
  • •
    \(\rho=\) specific resistance / resistivity
  • •
    \(\kappa=\dfrac{1}{\rho}\)
  • •
    \(\kappa=\dfrac{l}{aR}\)
  • •
    \(\dfrac{l}{a}=\) cell constant
  • •
    \(\dfrac{1}{R}=C\)
  • •
    \(\kappa=\dfrac{l}{a}\times C\)
  • •
    Specific conductance = conductance × cell constant
  • •
    Resistance = cell constant × resistivity
❖ Units:
  • •
    \(ohm^{-1}cm^{-1}\)
  • •
    \(S\,cm^{-1}\) [IOM 2007]
❖ Cell Constant: Determined with conductivity cell filled with \(N/50\) chemically pure KCl solution.
📚
EQUIVALENT AND MOLAR CONDUCTIVITY
▢ Equivalent Conductivity \((\lambda)\):
❖ Definition: Conductance of all ions produced from one gram equivalent of electrolyte dissolved in \(V\,cc\) solution when electrodes are 1 cm apart and whole solution is between them.
❖ Formulae:
  • •
    \(\lambda=V\kappa\)
  • •
    If concentration = \(N\) gram equivalent per litre: \(\lambda=\kappa\times\dfrac{1000}{N}\)
❖ Units: \(ohm^{-1}cm^2\,g\,eq^{-1}\) or \(S\,cm^2\,g\,eq^{-1}\)
❖ Dilution Effect: Increases with dilution.
▢ Molar Conductivity \((\mu\ or\ \Lambda_m)\):
❖ Definition: Conductance of all ions produced by one gram mole of electrolyte present in \(V\,cc\) solution when electrodes are 1 cm apart and solution lies between them.
❖ Formulae:
  • •
    \(\mu=\kappa V\)
  • •
    If concentration = \(M\) mole per litre: \(\mu=\kappa\times\dfrac{1000}{M}\)
❖ Units: \(ohm^{-1}cm^2mol^{-1}\) or \(S\,cm^2mol^{-1}\)
❖ Dilution Effect: Increases with dilution.
Relation of \(\mu\) and \(\lambda\)
  • •
    For 1:1 electrolyte, \(\mu=\lambda\), e.g. \(NaCl\), \(KBr\), \(NaNO_3\).
  • •
    For 1:2 or 2:1 electrolyte, \(\mu=2\lambda\), e.g. \(CaCl_2\), \(Na_2CO_3\), \(MgSO_4\).
  • •
    \(\text{Equivalent conductivity of ion}=\dfrac{\text{Molar conductivity of ion}}{\text{Charge on ion}}\)
📚
EFFECT OF DILUTION ON CONDUCTANCE
▢ Specific Conductance \((\kappa)\): Decreases with dilution for weak and strong electrolytes because number of ions per unit volume decreases.
▢ Weak Electrolyte: \(\lambda\) and \(\mu\) increase with dilution because volume increase effect is greater than decrease in \(\kappa\).
▢ Strong Electrolyte:
  • •
    Nearly 100% ionized at all concentrations.
  • •
    On dilution, \(\lambda\) and \(\mu\) increase due to increased ionic speed caused by less interionic attraction.
  • •
    \(\lambda\) and \(\mu\) do not change considerably on dilution.
📚
FACTORS AFFECTING ELECTROLYTIC CONDUCTANCE

Table 1: Factors

Factor
Effect
Solute-solute / interionic interaction
Greater interionic interaction → lower conductance
Solute-solvent interaction / hydration
Greater hydration → lower conductance
Solvent-solvent interaction / viscosity
More viscous medium → lower conductance
Temperature
\(\kappa\), \(\lambda\), \(\mu\) increase with temperature
▢ Hydration and Size:
  • •
    Ionic size order: \(Li^+
  • •
    Aqueous hydrated size order: \(Li^+(aq)>Na^+(aq)>K^+(aq)>Rb^+(aq)>Cs^+(aq)\)
  • •
    Ionic conductance maximum for Cs, minimum for Li in aqueous medium.
📚
SPECIAL CONDUCTANCE POINTS
▢ Oxidising and Reducing Power:
  • •
    Greater reducing power of metal → weaker oxidising action of cation.
  • •
    Li = most powerful reducing agent in aqueous solution; \(Li^+\) = weakest oxidising agent.
  • •
    Non-metals are oxidising agents; their anions act as reducing agents.
  • •
    Greater oxidising power of non-metal → weaker reducing action of its anion.
  • •
    Fluorine = most powerful oxidising agent; \(F^-\) = weakest reducing agent.
▢ Solubility Relation:
  • •
    \(\text{Solubility of sparingly soluble salt}=\dfrac{K\times1000}{\mu^\circ}\)
  • •
    \(\mu^\circ=\) molar conductivity at infinite dilution
▢ Debye-Huckel Equation:
  • •
    \(\mu=\mu^\circ-b\sqrt{C}\)
  • •
    \(b=\) Debye-Huckel constant
📚
KOHLRAUSCH'S LAW
📖
_*c
📝
Statement
Equivalent conductance of electrolyte at infinite dilution is sum of two values, one depending on cation and another on anion.
📝
Formula
\(\lambda*\infty=\lambda_c+\lambda_a\)
📝
Important Points
  • •
    Valid at any dilution but applied only at infinite dilution.
  • •
    Valid for both strong and weak electrolytes.
  • •
    At infinite dilution, attraction among dissimilar ions and repulsion among similar ions is approximately nil.
  • •
    Used to determine molar conductivity of weak electrolyte at infinite dilution.
📚
ELECTRODE POTENTIAL
▢ Definition: Potential difference developed between metal and solution when metal is placed in solution of its ions.
▢ Example: Zinc plate in \(Zn^{2+}\) solution becomes negatively charged and develops electrode potential.
▢ Oxidation Potential:
  • •
    Potential associated with oxidation reaction.
  • •
    Electrode acts as anode.
  • •
    \(M\rightarrow M^{n+}+ne^-\)
▢ Reduction Potential:
    ◉ **type: bullet
  1. Potential associated with reduction reaction.
  2. Electrode acts as cathode.
  3. \(M^{n+}+ne^-\rightarrow M\)
▢ **note:
    ▢ Depends On:
    • •
      Nature of metal
    • •
      Concentration of metallic ion in solution
    • •
      Temperature of solution
    📚
    STANDARD ELECTRODE POTENTIAL
    ▢ Definition: Potential difference in volts developed in a cell consisting of two electrodes: pure metal in contact with molar solution of its ion and normal hydrogen electrode at \(25^\circ C\).
    ▢ Important Points:
    • •
      Single electrode potential cannot be measured accurately alone.
    • •
      Potential difference between two electrodes can be measured.
    • •
      Normal hydrogen electrode / Standard hydrogen electrode = standard reference electrode.
    • •
      Electrode potential of NHE/SHE assigned zero volt.
    • •
      SHE difficult to prepare and maintain; replaced by secondary reference electrodes.
    • •
      Secondary reference electrodes: calomel electrode, silver-silver chloride electrode.
    • •
      Standard electrode potential measured by voltmeter.
    • •
      Positive electrode potential means species gets reduced with SHE.
    • •
      Standard oxidation potential = negative of standard reduction potential.
    ▢ Electrochemical Series: Arrangement of elements in increasing standard reduction potential.
    📚
    ELECTROCHEMICAL SERIES

    Table 1: Electrochemical series

    Electrode
    Half-cell reduction reaction
    \(E^\circ\) V
    \(Li^+/Li\)
    \(Li^+ + e^- \rightarrow Li\)
    -3.045
    \(K^+/K\)
    \(K^+ + e^- \rightarrow K\)
    -2.925
    \(Ca^{2+}/Ca\)
    \(Ca^{2+}+2e^-\rightarrow Ca\)
    -2.866
    \(Na^+/Na\)
    \(Na^+ + e^-\rightarrow Na\)
    -2.714
    \(Mg^{2+}/Mg\)
    \(Mg^{2+}+2e^-\rightarrow Mg\)
    -2.363
    \(Al^{3+}/Al\)
    \(Al^{3+}+3e^-\rightarrow Al\)
    -1.662
    \(Zn^{2+}/Zn\)
    \(Zn^{2+}+2e^-\rightarrow Zn\)
    -0.763
    \(Fe^{2+}/Fe\)
    \(Fe^{2+}+2e^-\rightarrow Fe\)
    -0.440
    \(Sn^{2+}/Sn\)
    \(Sn^{2+}+2e^-\rightarrow Sn\)
    -0.136
    \(Pb^{2+}/Pb\)
    \(Pb^{2+}+2e^-\rightarrow Pb\)
    -0.126
    \(H^+/H_2,Pt\)
    \(2H^+ +2e^-\rightarrow H_2\)
    0.000
    \(Cu^{2+}/Cu\)
    \(Cu^{2+}+2e^-\rightarrow Cu\)
    +0.340
    \(Ag^+/Ag\)
    \(Ag^+ + e^-\rightarrow Ag\)
    +0.800
    \(Cl^-/Cl_2\)
    \(Cl_2+2e^-\rightarrow2Cl^-\)
    +1.360
    \(Au^{3+}/Au\)
    \(Au^{3+}+3e^-\rightarrow Au\)
    +1.500
    \(F^-/F_2,Pt\)
    \(F_2+2e^-\rightarrow2F^-\)
    +2.870
    Trend
    • •
      Top metals: strongest reducing agents, weakest oxidising agents.
    • •
      Bottom species: strongest oxidising agents, weakest reducing agents.
    • •
      Standard oxidation potentials have same numerical values but opposite signs.
    📚
    CHARACTERISTICS OF ELECTROCHEMICAL SERIES
    ▢ Metal Activity:
    • •
      Top metals are strongly electropositive / weakly electronegative.
    • •
      Metal with lower reduction potential displaces metal with higher reduction potential from salt solution.
    • •
      \(CuSO_4\) cannot be stored in iron bottle.
    • •
      \(AgNO_3\) cannot be stored in copper bottle.
    • •
      Metals above hydrogen are easily rusted compared with metals below hydrogen.
    • •
      Metals above hydrogen displace hydrogen from dilute acids.
    • •
      K, Na, Ca displace hydrogen from water.
    • •
      Iron and metals above it decompose steam to liberate \(H_2\).
    • •
      Oxides of Hg and metals below hydrogen decompose on heating.
    ▢ Hydroxides and Salts:
    • •
      Hydroxides of metals in upper part are strongly basic; their salts do not undergo hydrolysis.
    • •
      Hydroxides of metals in lower part are weakly basic; their salts undergo hydrolysis.
    ▢ Non-metals:
      ◉ _*type: bullet
    1. Substances stronger oxidising than hydrogen are placed below hydrogen.
    2. Non-metal lower in series with high reduction potential displaces another non-metal with lower reduction potential.
    ▢ Reducing Power:
      ◉ **type: bullet
    1. Tendency to donate electron \(\propto E^\circ*{oxid}\).
    2. Tendency to donate electron \(\propto \dfrac{1}{E^\circ*{red}}\).
    3. If \(E^\circ*{red}\) of metallic cations x, y, z are \(0.52\), \(-3.03\), \(-1.18\) V, reducing power of metals: \(y>z>x\).
    📚
    INFORMATION FROM STANDARD REDUCTION POTENTIAL
    📖
    _*c
    📝
    Key Points
      ■ **type: bullet
    1. Higher positive \(E^\circ*{red}\) → lower reducing strength → higher oxidising strength.
    2. Li has least reduction potential \((-3.045V)\) → strongest reducing agent.
    3. \(F_2\) has maximum reduction potential \((+2.870V)\) → strongest oxidising agent / weakest reducing agent.
    4. More active metal \((E^\circ*{red}\ low, E^\circ*{oxid}\ high)\) displaces less active metal from salt solution.
    5. CuSO4 is not stored in Al container because Cu gets reduced by Al.
    6. Metals above \(H_2\) in activity series displace hydrogen from hydracids and get oxidized.
    Concentration and temperature effect
    • •
      Reduction potential of reducing agent decreases with rise in temperature and concentration of reducing solution.
    • •
      Reduction potential of oxidising agent increases with rise in temperature and concentration of oxidising solution.
    📚
    ELECTROLYTIC CELL AND ELECTROCHEMICAL CELL
    ▢ Electrolytic Cell:
    • •
      Electrical energy converted into chemical energy.
    • •
      Used for purification of metals.
    • •
      Electrons flow from cathode to anode through internal supply. [MOE]
    ▢ Electrochemical / Galvanic / Voltaic Cell:
    • •
      Chemical energy converted into electrical energy.
    • •
      Examples: dry cell, lead storage battery, Daniel cell.

    Table 1: Signs and reactions

    Cell
    Anode sign
    Cathode sign
    Anode reaction
    Cathode reaction
    Electrolytic cell
    +
    -
    Oxidation
    Reduction
    Galvanic cell
    -
    +
    Oxidation
    Reduction

    Table 2: Difference between electrolytic and galvanic cells

    Feature
    Electrolytic cell
    Galvanic cell
    Energy conversion
    Electrical → chemical
    Chemical → electrical
    Electrode signs
    Anode +, cathode -
    Anode -, cathode +
    Ion discharge
    Ions discharged on both electrodes
    Ions discharged only on cathode
    If electrodes inert
    Electrolyte concentration decreases during current flow
    Anodic half-cell concentration increases, cathodic half-cell concentration decreases
    Compartments
    Both electrodes may be in same compartment
    Electrodes in different compartments
    Salt bridge
    Not used
    Used
    📚
    DANIEL CELL
    ▢ Components:
    • •
      Zn electrode dipping in \(1M\ ZnSO_4\)
    • •
      Cu electrode dipping in \(1M\ CuSO_4\)
    ▢ Electrodes:
    • •
      Zn electrode = anode = negative electrode; oxidation occurs [MOE 2061]
    • •
      Cu electrode = cathode = positive electrode; reduction occurs
    • •
      Cu has more reduction potential than Zn
    ▢ Half Reactions:
    • •
      Anode: \(Zn\rightarrow Zn^{2+}+2e^-\)
    • •
      Cathode: \(Cu^{2+}+2e^-\rightarrow Cu\)
    • •
      Cell reaction: \(Zn+Cu^{2+}\rightarrow Zn^{2+}+Cu\)
    ▢ Electron and Current Flow:
    • •
      Electrons flow externally from Zn electrode to Cu electrode.
    • •
      Current flows outside cell from Cu to Zn, i.e. cathode to anode.
    ▢ Notation:
    • •
      \(Zn/Zn^{2+}(1M)\ ||\ Cu^{2+}(1M)/Cu\)
    • •
      LHS = oxidation half-cell = anode = negative
    • •
      RHS = reduction half-cell = cathode = positive
    📚
    SALT BRIDGE
    📖
    **c
    📝
    Definition
    Glass tube filled with gel-like agar-agar or gelatin mixed with electrolyte such as KCl, \(KNO_3\), \(NH_4NO_3\).
    📝
    Functions
      ■ **type: bullet
    1. Completes electrical circuit.
    2. Maintains electrical neutrality in both compartments.
    3. Avoids liquid junction potential.
    📝
    Electrolyte Selection
      ■ **type: bullet
    1. Electrolytes with nearly same cation and anion mobility used.
    2. Examples: KCl, \(KNO_3\), \(NH_4NO_3\), \(NH_4Cl\).
    3. KCl salt bridge not used with Ag or Pb salts because AgCl and PbCl2 are insoluble.
    📝
    **note
      ■ title: Inert electrode
      ■ data: If one constituent is gas or both oxidized and reduced forms are in solution, electrode is set up by immersing Pt wire, e.g. \(Pt/H_2(1atm)/H^+||Cu^{2+}/Cu\).
    📚
    EMF OF CELL
    📖
    **c
    📝
    Formulae
      ■ **type: bullet
    1. \(E*{cell}=E*{electrode}(RHS)-E*{electrode}(LHS)\)
    2. \(E*{cell}=E*{red}(RHS)-E*{red}(LHS)\)
    3. \(E*{cell}=E*{cathode,red}-E*{anode,red}\)
    4. \(E*{cell}=E*{oxid}(LHS)+E*{red}(RHS)\)
    📝
    _*note
      📚
      NERNST EQUATION
      📖
      **c
      📝
      For Metal Electrode
      📄
      Reaction
      \(M^{n+}+ne^-\rightarrow M\)
      📄
      Equation
      \(E=E^\circ+\dfrac{RT}{nF}\ln\dfrac{[M^{n+}]}{[M]}\)
      📄
      At 298 K
      \(E=E^\circ+\dfrac{0.0591}{n}\log[M^{n+}]\)
      📄
      Reason
      For metal-metal ion electrode, \([M]=1\) because active mass of solid is unity.
      📝
      Terms
        ■ **type: bullet
      1. \(E=\) electrode potential
      2. \(E^\circ=\) standard electrode potential
      3. \(R=8.314\,J\,mol^{-1}K^{-1}\)
      4. \(T=\) temperature in kelvin
      5. \(n=\) electrons in half-cell reaction
      6. \(F=96500\,C\)
      7. \([M]=\) activity of metal = unity for pure metal
      8. \([M^{n+}]=\) molarity of oxidized form
      📝
      General At 298 K
      \(E=E^\circ-\dfrac{0.0591}{n}\log\dfrac{[product]}{[reactant]}\)
      📝
      Cell Nernst Equation
        ■ **type: bullet
      1. \(E*{cell}=E^\circ*{cell}+\dfrac{0.0591}{n}\log\dfrac{[Ox.\ of\ cathode][Red.\ of\ anode]}{[Ox.\ of\ anode][Red.\ of\ cathode]}\)
      2. For Daniel cell: \(E*{cell}=E^\circ*{cell}+\dfrac{0.0591}{2}\log\dfrac{[Cu^{2+}][Zn]}{[Zn^{2+}][Cu]}\)
      3. Since \([Zn]=[Cu]=1\): \(E*{cell}=E^\circ*{cell}+\dfrac{0.0591}{2}\log\dfrac{[Cu^{2+}]}{[Zn^{2+}]}\)
      📝
      Relation with Equilibrium Constant
      \(E^\circ*{cell}=\dfrac{0.0591}{n}\log K*{eq}\)
      📚
      ELECTRODE POTENTIAL: NUMERICAL POINTS
      ▢ _*table:
        Dilution effect
        On dilution of electrolyte solution, reduction electrode potential decreases; oxidation electrode potential increases.
        📚
        COMMERCIAL CELLS
        ▢ Primary Cells:
        ❖ Characters:
        • •
          Definite life.
        • •
          Become dead with time.
        • •
          Cannot be recharged.
        ❖ Examples:
        1. Dry / Leclanche cell
        2. Mercury cell
        ❖ Leclanche Cell:
        • •
          Cylindrical zinc container = anode / negative electrode.
        • •
          Graphite rod = cathode / positive electrode.
        • •
          Electrolyte: mixture of \(MnO_2\), \(NH_4Cl\), \(ZnCl_2\).
        • •
          \(MnO_2\) = depolarizer.
        • •
          Uses: torches, toys, flashlights, calculators, tape recorders.
        ▢ Secondary Cells / Storage Cells / Accumulators:
        ❖ Characters:
        • •
          Rechargeable by passing direct current.
        • •
          Electrode reactions reversible.
        ❖ Examples:
        1. Lead storage battery
        2. Nickel-cadmium cell
        ❖ Lead Storage Battery:
        • •
          Anode: lead rod.
        • •
          Cathode: lead plates coated with \(PbO_2\).
        • •
          Electrolyte: aqueous \(H_2SO_4\).
        • •
          Most commonly used battery in automobiles.
        • •
          \(Pb+PbO_2+2H_2SO_4 \rightleftharpoons 2PbSO_4+2H_2O\)
        • •
          On charging: \(PbSO_4\) decomposed; \(PbO_2\) and \(H_2SO_4\) regenerated.
        • •
          On discharging: \(PbO_2\) and \(H_2SO_4\) consumed; \(PbSO_4\) deposited.
        ▢ Fuel Cell:
        • •
          Converts chemical energy of fuels directly into electrical energy.
        • •
          Example: \(H_2-O_2\) fuel cell.
        • •
          Theoretical efficiency = 100%.
        • •
          Practical efficiency = 70%.
        • •
          \(NaOH\) used as electrolyte.
        • •
          Efficient and pollution-free.
        • •
          Used in spacecrafts for electrical power.
        📚
        CORROSION
        ▢ Definition: Electrochemical process.
        ▢ Rust:
        • •
          Rust = \(Fe_2O_3\cdot xH_2O\)
        • •
          Rust is a mixture of \(Fe_2O_3\) and \(Fe(OH)_3\).
        ▢ Prevention:
        • •
          Barrier protection
        • •
          Sacrificial protection
        • •
          Electrical protection
        • •
          Galvanizing: covering iron with zinc
        • •
          Conductivity cell is platinized to avoid polarization effect
        📚
        READ & DIGEST
        ▢ Important Points:
          ◉ **type: bullet
        1. Ohm's law is followed by metallic and electrolytic conductors, but Faraday's law only by electrolytic conductors.
        2. During electrolysis, species discharged at cathode are cations.
        3. In electrolytic cell, electrons flow from cathode to anode through internal supply.
        4. In galvanic cell, electrons flow from anode to cathode through external circuit.
        5. In electroplating, article to be electroplated acts as cathode.
        6. Corrosion is an electrochemical process.
        7. A smuggler cannot cover gold with iron by depositing iron on gold surface because gold has higher standard reduction potential than iron.
        8. In Leclanche dry cell, anode is zinc container.
        9. Saturated \(KNO_3\) solution is used in salt bridge because velocities of \(K^+\) and \(NO_3^-\) are nearly same.
        10. Reference electrode is made from calomel \((Hg_2Cl_2)\).
        11. \(\kappa\), \(\lambda\), \(\mu\) increase with increase in temperature.
        12. Metals above hydrogen are easily rusted than those below.
        13. Electrolysis of dilute \(H_2SO_4\): \(H_2\) at cathode, \(O_2\) at anode.
        14. Electrolysis of concentrated \(H_2SO_4\): \(H_2S_2O_8\) / Marshall acid and \(H_2SO_5\) / Caro's acid at anode; \(H_2\) at cathode.
        15. Pure water does not conduct electricity because it is almost non-ionized.
        16. Solid \(NaCl\) is bad conductor; graphite is good conductor.
        17. Electrolysis of molten NaCl gives sodium and chlorine.
        18. Electrolysis of aqueous NaCl gives hydrogen and chlorine.
        19. If mercury cathode is used in electrolysis of NaCl solution, \(Na^+\) deposits at cathode forming Na-amalgam.
        20. Products of electrolysis of aqueous \(Na_2SO_4\) are \(H_2\) and \(O_2\).
        21. During electrolysis, if cathode is pulled out, ions start moving randomly.
        22. In aqueous dilute \(H_2SO_4\), ion discharged at anode is \(OH^-\), but gas evolved is \(O_2\).
        23. In NaCl electrolysis with Pt electrode, \(H_2\) liberated at cathode; with Hg cathode, Na-amalgam forms.
        24. Zn gives \(H_2\) with \(H_2SO_4\) and HCl but not with \(HNO_3\), because \(NO_3^-\) is reduced preferentially.
        25. Apparatus used to measure quantity of electricity = coulometer.
        26. Faraday's laws are related to equivalent weight of electrolytes.
        27. 1 coulomb = \(6.28\times10^{18}\) electrons.
        28. Total charge of one mole monovalent ion = \(9.65\times10^4C\).
        29. All galvanic cells do not contain porous plate.
        30. If salt bridge is removed, voltage drops to zero.
        31. Stronger oxidising agent → greater reduction potential.
        📚
        **MCQ
          1. A substance with initial concentration of 'a' \(mol\,dm^{-3}\) reacts according to zero order kinetics. The time it takes for completion of the reaction is \((K=\) rate constant\()\)
          2. \(K/a\)
          3. \(a/2K\)
          4. \(a/K\)
          5. \(2K/a\)
          6. c
          7. \(t*{completion}=a/K\) for zero order reaction.
            1. In Arrhenius plot, intercept is equal to
            2. \(-E_a/R\)
            3. \(\ln A\)
            4. \(\ln K\)
            5. \(\log*{10} a\)
            6. b
            7. Arrhenius plot of \(\ln K\) vs \(1/T\) has intercept \(\ln A\).
              1. Which of the following kind of catalysis can be explained by the adsorption theory?
              2. Homogeneous catalysis
              3. Acid-base catalysis
              4. Heterogeneous catalysis
              5. Enzyme catalysis
              6. c
              7. Adsorption theory explains heterogeneous catalysis.
                1. The time for 90% of the 1st order reaction to be completed is
                2. 1.1 times that of half-life
                3. 2.2 times that of half-life
                4. 10 times that of half-life
                5. 3.3 times that of half-life
                6. d
                7. For first order reaction, \(t*{90\%}=3.3t*{1/2}\).
                  1. In a reaction: \(2X+Y\rightarrow M+N\), the concentration of Y is kept constant and that of X is tripled. Now, the rate of reaction will
                  2. Increase 3 times
                  3. Decrease 3 times
                  4. Increase 6 times
                  5. Increase 9 times
                  6. d
                  7. For rate depending on \([X]^2\), tripling X increases rate by \(3^2=9\) times.
                    1. The order of a reaction with respect to \(OH^-\) is \(-1\). The \(OH^-\) species
                    2. Acts as catalyst
                    3. Acts as inhibitor
                    4. Will always be involved in the overall reaction
                    5. Helps in increasing reaction rate
                    6. b
                    7. Negative order means increasing that species decreases rate; it behaves as inhibitor.
                    Q1.
                    The number of Coulombs of electricity required to deposit 0.3 mole of \(Cu^{++}\) is
                    📅MOE Model
                    Q2.
                    The amount of electricity required to deposit 0.2 mole of \(Ag^+\) is
                    📅MOE 2008
                    Q3.
                    When sodium chloride is electrolysed, which of the following reaction takes place at anode?
                    📅MOE 2060
                    Q4.
                    Electrochemical equivalent refers to mass deposited in which of the following cases?
                    📅MOE 2060
                    Q5.
                    \(CuSO_4\) is electrolyzed by passing 1 Faraday of current. At anode, there is
                    📅MOE 2003
                    Q6.
                    In electrolysis, amount of silver deposited equivalent of 112 ml of hydrogen is
                    📅MOE 2054
                    Q7.
                    An electric current 0.25 Amp was passed through acidified water for 2 hours. What will be the volume of \(H_2\) produced at NTP? \((1\,Coulomb\) of electricity produced \(10^{-5}\,gm\) of \(H_2)\)
                    📅IOM 08
                    Q8.
                    Given that \(Zn^{++} + 2e^- \rightarrow Zn, E^\circ = -0.76V\) and \(Mg^{++} + 2e^- \rightarrow Mg, E^\circ = -2.37V\). When Zn dust is added to \(MgCl_2\) solution then
                    📅B.E. 2065
                    Q9.
                    In the electrolysis of water, 224 litres of pure hydrogen was liberated at NTP. The weight of oxygen liberated will be
                    📅MOE 2065
                    Q10.
                    The weight of Ag deposited when 465.5 coulomb electricity is passed is
                    📅I.E. 2009
                    Q11.
                    When \(CuSO_4\) solution is electrolysed between the two electrodes then at anode
                    Q12.
                    Electrolysis of \(H_2O\) produces 112 ml of \(O_2\) at NTP, then mass of hydrogen gas produced is
                    📅KU 2008
                    Q13.
                    The conductivity of strong electrolyte
                    📅BPKIHS 2007
                    Q14.
                    Three faradays of charge is supplied to bivalent metal. What is the number of electrons involved?
                    📅BPKIHS 1999
                    Q15.
                    How many faradays of charge is required to convert 1 mole of \(MnO_4^-\) to \(Mn^{2+}\)?
                    📅I.E.
                    Q16.
                    If one end of a piece of metal is heated, the other end becomes hot after some time. This is due to
                    Q17.
                    The unit of specific conductivity is
                    Q18.
                    The unit of equivalent conductivity is
                    Q19.
                    The best conductor of electricity is a 1 M solution of
                    Q20.
                    Which of the following is a poor conductor of electricity?
                    Q21.
                    The molar conductivity of a strong electrolyte
                    Q22.
                    In the electrolysis of NaCl
                    Q23.
                    On the electrolysis of aqueous solution of sodium sulphate, on cathode we get
                    Q24.
                    In the electrolytic cell, flow of electrons is from
                    Q25.
                    The metal that cannot displace hydrogen from dilute hydrochloric acid is
                    Q26.
                    Faraday's laws of electrolysis will fail when
                    Q27.
                    One faraday of electricity will liberate one gram atom of the metal from a solution of
                    Q28.
                    When one ampere current flows for 1 second through a conductor, this quantity of electricity is called
                    Q29.
                    The charge required to liberate 11.5 g of sodium is
                    Q30.
                    On passing 0.1 faraday of electricity through aluminium chloride, the amount of aluminium metal deposited on cathode is \((Al=27)\)
                    Q31.
                    If the electrodes in the cell are 2.2 cm apart and have an area of 3.8 sq. cm, then the cell constant is
                    Q32.
                    Pure water is poor conductor of electricity because it
                    Q33.
                    Out of Cu, Ag, Fe and Zn, the metal which can displace all others from their salt solutions is
                    Q34.
                    The number of electrons passing per second through a cross-section of copper wire carrying \(10^{-6}\) amperes of current per second is found to be
                    Q35.
                    Electrolytic cell is used to convert
                    Q36.
                    In galvanic cells
                    Q37.
                    As a lead storage battery is charged
                    Q38.
                    If a salt bridge is removed between the two half cells, the voltage
                    Q39.
                    The reference electrode is made from which of the following?
                    Q40.
                    The correct order of chemical reactivity with water according to electrochemical series is
                    Q41.
                    The standard reduction potential for \(Fe^{2+}/Fe\) and \(Sn^{2+}/Sn\) electrodes are -0.44 and -0.14 volt respectively. For the cell reaction: \(Fe^{2+}+Sn\rightarrow Fe+Sn^{2+}\), the standard emf is
                    Q42.
                    Prevention of corrosion of iron by Zn coating is called
                    Q43.
                    The number of coulombs required for the deposition of 107.870 g of silver is
                    Q44.
                    The standard reduction potential values of the three metallic cations X, Y and Z are 0.52, -3.03 and -1.18 V respectively. The order of reducing power of corresponding metals is
                    Q45.
                    \(Cu(s)+2Ag^+(aq)\rightarrow Cu^{2+}(aq)+2Ag(s)\), the reduction half cell reaction is
                    Q46.
                    What would be the weight of silver deposited on passing 965 coulombs of electricity in a solution of silver nitrate?
                    Q47.
                    The equivalent conductance at infinite dilution of a weak acid such as HF
                    Q48.
                    Cell reaction is spontaneous when
                    Q49.
                    In the electrochemical reaction \(2Fe^{3+}+Zn\rightarrow Zn^{2+}+2Fe^{2+}\), increasing the concentration of \(Fe^{3+}\)
                    Q50.
                    Basically, corrosion is
                    Q51.
                    Consider the following \(E^\circ\) values: \(E^\circ_{Fe^{3+}/Fe^{2+}}=+0.77V\), \(E^\circ_{Sn^{2+}/Sn}=-0.14V\). Under standard conditions, the potential for the reaction \(Sn(s)+2Fe^{3+}(aq)\rightarrow2Fe^{2+}(aq)+Sn^{2+}(aq)\) is
                    Q52.
                    The highest electrical conductivity of the following aqueous solutions is of
                    Q53.
                    Which of the following does not conduct electricity?
                    Q54.
                    \(E^\circ_{Cu}=0.34V\), \(E^\circ_{Zn}=-0.76V\). A Daniel cell contains 0.1 M \(ZnSO_4\) solution and 0.01 M \(CuSO_4\) solution at its electrodes. E.M.F of the cell is
                    Q55.
                    Electrochemical equivalent of an element is
                    Q56.
                    For the reactions: \(C+O_2\rightarrow CO_2, \Delta H=-393\,KJ\); \(2Zn+O_2\rightarrow2ZnO, \Delta H=-412\,KJ\)
                    Q57.
                    The amount of ion discharged during electrolysis is not directly proportional to
                    Q58.
                    Which of the following reactions is not possible?
                    Q59.
                    How many coulombs are required for the oxidation of 1 mol of \(H_2O_2\) to \(O_2\)?
                    Q60.
                    Which of the following cell convert chemical energy of \(H_2\) and \(O_2\) directly into electrical energy?
                    Q61.
                    In a hydrogen-oxygen fuel cell, combustion of hydrogen occurs to
                    Q62.
                    For spontaneity of a cell, which is correct?
                    Q63.
                    Which of the following statement is true for electrochemical Daniel cell?
                    Q64.
                    Which one of the following will increase the voltage of the cell? \(Sn(s)+2Ag^+(aq)\rightarrow Sn^{2+}(aq)+2Ag(s)\)
                    Q65.
                    When a piece of copper wire is immersed in a solution of \(AgNO_3\), the colour of the solution becomes blue. This is due to the
                    Q66.
                    Which of the following represents Faraday's 1st law?
                    Q67.
                    The molar conductivity of a strong electrolyte