Oxidation number of chromium in K₂Cr₂O₇ is
A +3
B +4
C +6
D +7
Let oxidation number of Cr = x
2(+1) + 2x + 7(−2) = 0
2 + 2x − 14 = 0 → 2x = 12 → x = +6
Oxidation number of nitrogen in NH₄⁺ ion is
A −1
B −2
C −3
D +1
Let N = x
x + 4(+1) = +1 → x + 4 = 1 → x = −3
Which element undergoes reduction in the reaction 2Al + Fe₂O₃ → 2Fe + Al₂O₃
A aluminium
B iron
C oxygen
D both iron and oxygen
Fe³⁺ in Fe₂O₃ gains electrons to form Fe⁰ → reduction.
In the same reaction, aluminium acts as
A oxidizing agent
B reducing agent
C catalyst
D electrolyte
Al is oxidized from 0 to +3, donating electrons.
Which species is simultaneously oxidized and reduced
A Cl₂
B SO₂
C H₂O₂
D Zn
In 2H₂O₂ → 2H₂O + O₂, oxygen undergoes both oxidation and reduction.
The balancing of redox reactions in basic medium requires addition of
A H⁺ only
B OH⁻ only
C H₂O only
D H⁺ and OH⁻ both
In basic medium, excess H⁺ is neutralized using OH⁻.
In acidic medium, Cr₂O₇²⁻ is reduced to Cr³⁺ by gaining
A 3 electrons
B 4 electrons
C 6 electrons
D 7 electrons
Each Cr changes from +6 to +3 → gain of 3 electrons × 2 = 6 electrons.
A concentration cell works due to difference in
A electrodes
B electrolytes
C concentration
D temperature
EMF arises from concentration gradient, not different electrodes.
In a concentration cell, emf becomes zero when
A electrodes are removed
B concentrations become equal
C temperature is zero
D salt bridge is broken
No concentration gradient → no driving force.
The cell Zn | Zn²⁺ (0.1 M) || Zn²⁺ (1.0 M) | Zn is
A galvanic cell
B electrolytic cell
C concentration cell
D fuel cell
Same electrodes, different concentrations.
In the above cell, electrons flow from
A higher concentration to lower concentration
B lower concentration to higher concentration
C cathode to anode
D salt bridge to electrode
Dilute solution acts as anode; electrons flow to concentrated side.
The emf of a concentration cell depends on
A E°cell
B difference in concentration
C nature of electrode
D size of electrode
E°cell = 0 for concentration cells; emf arises from concentration ratio.
Standard emf of a concentration cell is
A always positive
B always negative
C zero
D infinity
Same electrodes → E°cathode = E°anode.
A concentration cell converts
A chemical energy into electrical energy
B concentration gradient into electrical energy
C electrical energy into chemical energy
D heat into electrical energy
Work is obtained by equalizing concentration difference.
The emf of a Daniell cell at 25°C is given by E = 1.10 − (0.0591/2) log([Zn²⁺]/[Cu²⁺]). If [Zn²⁺] = [Cu²⁺], the emf will be
A 0.00 V
B 0.59 V
C 1.10 V
D 1.69 V
log(1) = 0 → E = E° = 1.10 V
If reaction quotient Q < 1, then cell emf
A is less than E°
B equals E°
C is greater than E°
D becomes zero
E = E° − (0.0591/n) log Q. If Q < 1 → log Q negative → E increases.
If Q > K, then reaction will
A proceed forward
B proceed backward
C stop permanently
D become explosive
System shifts to reduce Q and reach equilibrium.
For a spontaneous reaction at given conditions
A E > 0
B E = 0
C E < 0
D E = E° only
Positive emf indicates spontaneity.
The relation between equilibrium constant and standard emf is
A E° = (RT/F) ln K
B E° = (RT/nF) ln K
C E° = −(RT/nF) ln K
D E° = (nF/RT) ln K
Derived from ΔG° = −nFE° and ΔG° = −RT ln K.
Large value of equilibrium constant implies
A E° small
B reaction incomplete
C reaction strongly product favored
D ΔG° positive
Large K → products dominate → E° large and positive.
Conductivity (κ) of a solution decreases on dilution because
A number of ions per unit volume decreases
B degree of ionization decreases
C temperature decreases
D ion mobility decreases
Dilution spreads ions over larger volume.
Molar conductance (Λm) increases on dilution because
A ions disappear
B interionic attraction decreases
C solution evaporates
D solvent conductivity increases
Ions move more freely at larger separation.
The plot of Λm versus √c for strong electrolytes is
A straight line with positive slope
B straight line with negative slope
C curved line
D horizontal line
Λm decreases linearly with √c due to interionic effects.
For weak electrolytes, the plot of Λm versus √c is
A straight line
B horizontal line
C curved line
D zig-zag line
Degree of ionization changes significantly with dilution.
Limiting molar conductivity of weak electrolytes is calculated using
A Ostwald’s law
B Kohlrausch law
C Raoult’s law
D Faraday’s law
Λm°(weak) = Λm°(strong electrolytes) − ionic contributions.
Degree of dissociation of weak electrolyte increases with
A increase in concentration
B decrease in temperature
C dilution
D addition of common ion
According to Ostwald’s dilution law.
In a dry cell, the electrolyte is
A liquid NH₄Cl
B aqueous NH₄Cl
C moist paste of NH₄Cl and ZnCl₂
D solid NaCl
Dry cell uses paste electrolyte.
In dry cell, carbon rod acts as
A anode
B cathode
C electrolyte
D salt bridge
Carbon rod collects electrons; reduction occurs there.
In dry cell, zinc container acts as
A cathode
B anode
C electrolyte
D insulator
Zn undergoes oxidation.
Lead storage battery during discharge produces
A Pb and PbO₂
B PbSO₄ on both plates
C PbCl₂
D PbCO₃
Both electrodes convert to PbSO₄ during discharge.
During charging of lead storage battery
A PbSO₄ is formed
B Pb and PbO₂ are regenerated
C sulfuric acid is consumed
D battery acts as galvanic cell
Charging reverses discharge reaction.
Corrosion of iron is faster in coastal areas because of
A high temperature
B high humidity and salt
C low oxygen
D low moisture
Electrolytes enhance corrosion.
Rusting of iron is an example of
A purely oxidation
B purely reduction
C electrochemical corrosion
D physical corrosion
Anode and cathode reactions occur simultaneously.
The method of protecting iron by connecting it to a more reactive metal is called
A galvanization
B cathodic protection
C anodization
D alloying
Iron becomes cathode; reactive metal sacrifices.
Magnesium is preferred over zinc as sacrificial anode because
A Mg is cheaper
B Mg has lower reduction potential
C Mg is heavier
D Mg is noble metal
More negative E° → oxidizes more readily.
Corrosion can be prevented by
A painting
B oiling
C galvanization
D all of these
All methods isolate metal from air/moisture or provide sacrificial protection.
Fuel cells are preferred because they
A pollute environment
B have low efficiency
C directly convert chemical energy into electrical energy
D need frequent replacement
Fuel cells bypass heat engine losses.
In hydrogen–oxygen fuel cell, product formed is
A H₂O
B H₂O₂
C OH⁻ only
D H⁺ only
Hydrogen and oxygen combine to form water.
Fuel cells are used in spacecraft mainly because
A they are cheap
B they are light and efficient
C they do not need fuel
D they use nuclear reactions
High efficiency and water byproduct is useful.
The corrosion of iron in presence of copper is
A slower
B completely stopped
C faster
D unchanged
Cu is more noble; Fe becomes anode and corrodes faster.
Corrosion rate increases when
A pH increases
B pH decreases
C temperature decreases
D moisture decreases
Acidic medium accelerates corrosion.
Electroplating is carried out in
A galvanic cell
B electrolytic cell
C concentration cell
D fuel cell
External electricity drives non-spontaneous deposition.
During electroplating, the object to be plated is made
A anode
B cathode
C electrolyte
D salt bridge
Metal ions reduce and deposit on cathode.
Faraday’s first law of electrolysis relates mass deposited to
A voltage
B time only
C charge passed
D resistance
m ∝ Q.
One faraday of electricity corresponds to charge of approximately
A 965 C
B 9650 C
C 96500 C
D 9.65 × 10⁶ C
Charge on 1 mole of electrons.
During corrosion, electrons flow through
A electrolyte
B metal
C air
D salt bridge
Electron flow occurs through metal body.
During corrosion, ions flow through
A metal
B electrolyte/moisture
C air
D wire
Ionic conduction occurs through moisture layer.
Corrosion inhibitor works by
A increasing oxidation
B forming protective film
C increasing conductivity
D removing electrons
Film blocks electrochemical reactions.
In cathodic protection, protected metal acts as
A anode
B cathode
C electrolyte
D oxidizing agent
Cathode does not oxidize.
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.