Corrosion of iron is an example of
A spontaneous redox reaction
B non-spontaneous reaction
C electrolytic reaction only
D neutralization reaction
Corrosion occurs naturally without external energy.
Oxidation number of sulfur in Na₂S₂O₃ is
A +2
B +4
C +6
D average +2
In thiosulfate, the two sulfur atoms are not equivalent (one is −2 and one is +6), so average oxidation state is +2.
Oxidation number of nitrogen in NO₃⁻ is
A +3
B +4
C +5
D +6
Let N = x: x + 3(−2) = −1 → x − 6 = −1 → x = +5.
In acidic medium, permanganate ion (MnO₄⁻) is reduced to
A MnO₂
B Mn²⁺
C Mn³⁺
D MnO₄²⁻
In acidic medium, MnO₄⁻ reduces to Mn²⁺ (purple to colorless).
In neutral or alkaline medium, MnO₄⁻ is commonly reduced to
A Mn²⁺
B MnO₂
C Mn³⁺
D MnO₄²⁻ only
In neutral/alkaline medium, MnO₄⁻ reduces to MnO₂ (brown precipitate).
In acidic medium, dichromate ion (Cr₂O₇²⁻) is reduced to
A CrO₄²⁻
B Cr³⁺
C Cr²⁺
D Cr⁶⁺
Cr(VI) in dichromate is reduced to Cr(III) in acidic solution.
In the oxidation number method, balancing is based on
A balancing atoms first always
B balancing charges first always
C equalizing total increase and decrease in oxidation numbers
D adding H₂O only
Total oxidation number increase = total decrease for electron balance.
In a redox reaction, total electrons lost are
A more than electrons gained
B less than electrons gained
C equal to electrons gained
D sometimes unequal
Conservation of charge requires electrons lost = electrons gained.
A species with higher standard reduction potential is more likely to
A undergo oxidation
B act as reducing agent
C undergo reduction
D act as anode always
Higher E°(reduction) means greater tendency to gain electrons.
Which metal is strongest reducing agent based on standard reduction potentials
A Ag
B Cu
C Zn
D Au
More negative E° indicates stronger reducing tendency; Zn has much lower E° than Cu/Ag/Au.
In a galvanic cell, the anode is
A positive electrode
B negative electrode
C always platinum
D always copper
In galvanic cell, anode releases electrons → negative.
In an electrolytic cell, the anode is
A negative electrode
B positive electrode
C always zinc
D always cathode
In electrolytic cell, anode is connected to positive terminal of power supply.
Which statement is correct about a salt bridge
A electrons flow through salt bridge
B ions flow through salt bridge
C salt bridge increases oxidation number
D salt bridge acts as metal wire
Salt bridge provides ionic conduction; electrons flow through external wire.
If E°cell is negative, the reaction under standard conditions is
A spontaneous
B non-spontaneous
C at equilibrium
D impossible to reverse
Negative E° means ΔG° is positive; reaction won’t proceed spontaneously.
For a galvanic cell, if E°cell = 0.00 V, then K is
A 0
B 1
C 10
D 100
ΔG° = 0 → K = 1.
At 25°C, for a reaction involving 2 electrons, (0.0591/n) equals
A 0.1182
B 0.0591
C 0.02955
D 0.0148
0.0591/2 = 0.02955.
For a cell, E = E° − (0.0591/n) log Q. If Q = 1, then E equals
A 0
B E°
C −E°
D depends only on n
log(1)=0, so E = E°.
For Daniell cell, if [Zn²⁺] increases while [Cu²⁺] remains constant, cell emf
A increases
B decreases
C becomes zero
D becomes infinite
Q = [Zn²⁺]/[Cu²⁺] increases → logQ increases → E decreases.
For a hydrogen electrode at 25°C, E = 0.00 − 0.0591 pH. If pH = 2, E is
A −0.1182 V
B −0.0591 V
C +0.1182 V
D 0.00 V
E = −0.0591×2 = −0.1182 V.
If E°cell = 0.1182 V at 25°C and n = 2, then logK is
A 1
B 2
C 4
D 0.5
logK = (2×0.1182)/0.0591 = 4.
If K is very large, then E°cell is
A negative and large magnitude
B positive and large magnitude
C zero
D unrelated to K
Large K implies strongly product-favored, hence E° is positive.
Specific conductivity κ depends on
A number of ions per unit volume
B mobility of ions
C temperature
D all of these
κ depends on ion concentration, mobility and temperature.
Molar conductance Λm is related to κ and concentration (c in mol/L) by
A Λm = κ × c
B Λm = κ / c
C Λm = 1000 κ / c
D Λm = c / κ
Λm (S cm² mol⁻¹) = κ (S cm⁻¹) × 1000 / c.
If κ = 2×10⁻³ S cm⁻¹ and c = 0.01 mol/L, molar conductance is
A 0.2
B 20
C 200
D 2
Λm = 1000κ/c = 1000×2×10⁻³ / 0.01 = 200 S cm² mol⁻¹.
Weak electrolyte shows sharp increase in Λm on dilution because
A viscosity increases
B degree of ionization increases
C ions become heavier
D κ increases sharply always
Weak electrolytes ionize more on dilution.
Strong electrolyte shows only small increase in Λm on dilution because
A it does not ionize
B it is already almost completely ionized
C it forms precipitate
D it forms micelles
Dilution mainly increases mobility slightly.
In lead storage battery, during discharge, concentration of H₂SO₄
A increases
B decreases
C remains constant
D becomes zero immediately
H₂SO₄ is consumed forming PbSO₄ and water; acid density decreases.
A fully charged lead storage battery has acid with
A higher density
B lower density
C zero density
D same density always
Higher H₂SO₄ concentration → higher density indicates charge state.
In a dry cell, depolarizer used is
A H₂SO₄
B MnO₂
C NaOH
D CuSO₄
MnO₂ removes H₂ buildup and prevents polarization.
Nickel–cadmium battery is
A primary
B secondary
C fuel cell
D concentration cell
Ni–Cd is rechargeable.
Corrosion of iron is minimized by applying a coating of
A copper (thick)
B tin (but scratched)
C zinc (galvanization)
D mercury
Zinc provides sacrificial protection even if coating is scratched.
Tin coating on iron, if scratched, makes iron corrosion
A slower than usual
B faster than usual
C impossible
D unchanged
Tin is more noble; iron becomes anode at exposed area → rapid galvanic corrosion.
Zinc coating on iron, if scratched, still protects iron because
A Zn is more noble
B Zn acts as sacrificial anode
C Zn reduces oxygen
D Zn increases resistance
Zn oxidizes preferentially, protecting Fe.
Corrosion is fastest when iron is connected to
A magnesium
B zinc
C copper
D aluminium
Copper is more noble; iron becomes anode and corrodes quickly.
In electrochemical corrosion, the anode region is the place where
A oxygen is reduced
B metal dissolves as ions
C hydrogen is formed only
D no reaction occurs
Oxidation occurs at anode: M → Mⁿ⁺ + ne⁻.
The cathodic reaction in neutral water rusting commonly produces
A H⁺ ions
B OH⁻ ions
C Cl⁻ ions
D SO₄²⁻ ions
O₂ + 2H₂O + 4e⁻ → 4OH⁻ at cathode.
Pitting corrosion is especially dangerous because
A it forms uniform rust layer
B it is slow and visible
C it creates deep localized holes
D it increases metal thickness
Small pits can cause sudden failure even with little overall mass loss.
Rusting is faster in saline water mainly because
A salt decreases oxygen solubility
B salt acts as electrolyte increasing conductivity
C salt removes moisture
D salt forms protective film
Higher ionic conductivity accelerates electrochemical reactions.
A corrosion inhibitor generally works by
A increasing oxidation rate
B forming a barrier film on metal surface
C increasing electrolyte concentration
D increasing temperature
Inhibitors reduce contact and slow electrochemical reactions.
Cathodic protection is used mainly for
A plastics
B pipelines and ship hulls
C glassware
D rubber
Long metallic structures are protected by making them cathode.
The standard potential of SHE is fixed as 0 V because
A it is easiest to prepare
B it is most stable reference electrode
C it gives maximum emf always
D it prevents corrosion
SHE provides a standard reference point for electrode potentials.
If Ecell becomes zero during operation, the cell reaction has
A become irreversible
B reached equilibrium
C become faster
D become explosive
E = 0 implies no driving force; equilibrium achieved.
If a cell has E° > 0 but E < 0 under given conditions, the reaction is
A spontaneous forward
B spontaneous backward
C impossible both ways
D always at equilibrium
Actual E decides direction; negative E means reverse direction is spontaneous.
A metal with very negative E° value is generally
A noble metal
B strong oxidizing agent
C strong reducing agent
D inert metal
It readily loses electrons (oxidizes), reducing other species.
Electrochemical series is helpful in predicting
A color of ions
B feasibility of redox reaction
C boiling point
D solubility only
E° values help determine spontaneity and displacement reactions.
In a galvanic cell, oxidation is always written at
A cathode side
B anode side
C salt bridge side
D right side only
By definition, oxidation occurs at anode.
During charging of a secondary battery, the system acts as
A galvanic cell
B electrolytic cell
C concentration cell
D fuel cell
External current drives reverse (non-spontaneous) reaction.
For an electrolytic cell, cathode is the electrode where
A oxidation occurs
B reduction occurs
C electrons are produced
D anions are discharged only
Cathode is always reduction electrode in both galvanic and electrolytic cells.
In electrolysis of molten NaCl, the product at cathode is
A Cl₂
B Na
C NaCl
D NaOH
Na⁺ gains electrons: Na⁺ + e⁻ → Na at cathode.
In electrolysis of molten NaCl, the product at anode is
A Na
B H₂
C Cl₂
D O₂
Cl⁻ is oxidized: 2Cl⁻ → Cl₂ + 2e⁻.
Faraday’s second law states that mass deposited is proportional to
A current only
B time only
C equivalent weight of substance for same charge
D resistance only
For same quantity of electricity, masses deposited are proportional to their equivalent weights.