A redox reaction is one in which
A only oxidation occurs
B only reduction occurs
C oxidation and reduction occur simultaneously
D atoms are neither oxidized nor reduced
Oxidation and reduction are complementary; one cannot occur without the other.
Oxidation is defined as
A gain of electrons
B loss of electrons
C gain of protons
D decrease in oxidation number
Oxidation involves loss of electrons (OIL: Oxidation Is Loss).
Reduction is defined as
A loss of electrons
B loss of oxygen
C gain of electrons
D increase in oxidation number
Reduction involves gain of electrons (RIG: Reduction Is Gain).
The substance that undergoes oxidation acts as
A oxidizing agent
B reducing agent
C catalyst
D electrolyte
A reducing agent donates electrons and gets oxidized.
The substance that undergoes reduction acts as
A reducing agent
B oxidizing agent
C solvent
D buffer
Oxidizing agent accepts electrons and gets reduced.
In the reaction Zn + Cu²⁺ → Zn²⁺ + Cu, Zn acts as
A oxidizing agent
B reducing agent
C catalyst
D electrolyte
Zn loses electrons (oxidized) and reduces Cu²⁺.
In the same reaction, Cu²⁺ acts as
A reducing agent
B oxidizing agent
C catalyst
D solvent
Cu²⁺ gains electrons (reduced).
Oxidation number of oxygen in most of its compounds is
A −1
B −2
C 0
D +2
Oxygen generally shows −2 oxidation state, except in peroxides and OF₂.
Oxidation number of hydrogen in metal hydrides is
A +1
B −1
C 0
D +2
In metal hydrides (e.g., NaH), hydrogen exists as H⁻.
Oxidation number of an element in its elemental form is
A +1
B −1
C 0
D variable
Free elements always have oxidation number zero.
Which species is oxidized in the reaction
Fe²⁺ → Fe³⁺ + e⁻
A Fe³⁺
B Fe²⁺
C electron
D none
Fe²⁺ loses an electron → oxidation.
Which species is reduced in the reaction
Cl₂ + 2e⁻ → 2Cl⁻
A Cl⁻
B electrons
C Cl₂
D none
Chlorine gains electrons → reduction.
A reaction in which the same element is both oxidized and reduced is called
A redox reaction
B neutralization
C disproportionation
D combination
In disproportionation, one species undergoes both oxidation and reduction.
Which reaction is an example of disproportionation
A Zn + CuSO₄ → ZnSO₄ + Cu
B 2H₂ + O₂ → 2H₂O
C 2H₂O₂ → 2H₂O + O₂
D NaOH + HCl → NaCl + H₂O
Oxygen in H₂O₂ is both oxidized and reduced.
The method commonly used to balance redox reactions in acidic medium is
A hit and trial method
B oxidation number method
C ion–electron method
D algebraic method
Ion–electron (half-reaction) method is systematic and preferred.
An electrochemical cell converts
A heat energy into electrical energy
B chemical energy into electrical energy
C mechanical energy into electrical energy
D electrical energy into chemical energy
Redox reactions drive electron flow, producing electricity.
In a galvanic cell, oxidation occurs at
A cathode
B anode
C salt bridge
D electrolyte
Oxidation always occurs at the anode (An Ox).
In a galvanic cell, reduction occurs at
A anode
B cathode
C salt bridge
D wire
Reduction always occurs at the cathode (Red Cat).
Electrons flow in the external circuit from
A cathode to anode
B anode to cathode
C salt bridge to electrodes
D electrolyte to wire
Electrons are released at anode and consumed at cathode.
The standard electrode potential is measured relative to
A copper electrode
B zinc electrode
C standard hydrogen electrode
D silver electrode
SHE is assigned zero potential by convention.
The standard potential of hydrogen electrode is
A +1.00 V
B −1.00 V
C 0.00 V
D +0.76 V
By definition, E°(SHE) = 0 V.
The emf of a cell is given by
A E°cell = E°anode − E°cathode
B E°cell = E°cathode − E°anode
C E°cell = E°anode + E°cathode
D E°cell = 0
Cell emf is the potential difference driving electrons.
A positive value of E°cell indicates that the reaction is
A non-spontaneous
B spontaneous
C at equilibrium
D impossible
Positive emf means ΔG < 0 → spontaneous reaction.
The unit of electrode potential is
A ampere
B coulomb
C volt
D ohm
Potential difference is measured in volts.
Which of the following is a galvanic cell
A electrolytic cell
B Daniell cell
C electrolyzer
D electroplating cell
Daniell cell is a classic galvanic cell.
The Nernst equation relates electrode potential to
A temperature only
B concentration only
C reaction quotient and temperature
D pressure only
Nernst equation: E = E° − (RT/nF) ln Q.
At 25°C, the Nernst equation is written as
A E = E° − (0.0591/n) log Q
B E = E° + (0.0591/n) log Q
C E = E° − (RT/F) log Q
D E = E° − log Q
This is the simplified form at 298 K.
When Q = K, the cell potential becomes
A maximum
B minimum
C zero
D infinite
At equilibrium, ΔG = 0 → E = 0.
Increasing concentration of reactants generally
A decreases cell potential
B increases cell potential
C has no effect
D makes E = 0
Reaction quotient decreases, increasing E.
The number of electrons transferred in a redox reaction is denoted by
A F
B n
C R
D Q
n represents moles of electrons exchanged.
Electrical conductance of a solution depends on
A nature of electrolyte
B concentration
C temperature
D all of these
All these factors affect ion mobility and number.
Specific conductance decreases with dilution because
A number of ions decreases per unit volume
B ion mobility decreases
C temperature decreases
D viscosity increases
Dilution reduces ions per cm³.
Molar conductance increases with dilution because
A number of ions decreases
B ion mobility increases
C viscosity increases
D solvent evaporates
Reduced interionic attraction enhances mobility.
Unit of specific conductance is
A ohm
B ohm⁻¹ cm⁻¹
C ohm cm
D volt
Conductance per unit length and area.
Unit of molar conductance is
A ohm⁻¹
B ohm⁻¹ cm² mol⁻¹
C ohm cm⁻¹
D volt mol⁻¹
Accounts for conductance of one mole of electrolyte.
A primary battery cannot be
A recharged
B used once
C portable
D small
Primary batteries are irreversible.
Dry cell is an example of
A secondary battery
B fuel cell
C primary battery
D electrolytic cell
Dry cell is non-rechargeable.
Lead storage battery is an example of
A primary battery
B secondary battery
C fuel cell
D galvanometer
It can be recharged many times.
Corrosion of iron is an example of
A reduction
B oxidation
C physical change
D neutralization
Iron loses electrons to form rust.
Rust is mainly
A FeO
B Fe₂O₃
C hydrated Fe₂O₃
D Fe₃O₄
Rust is hydrated ferric oxide.
Corrosion is faster in
A dry air
B moist air
C vacuum
D inert gas
Moisture facilitates electrochemical reactions.
Galvanization protects iron by
A oxidation
B coating with zinc
C coating with copper
D alloying with carbon
Zinc acts as sacrificial anode.
Sacrificial protection works because the coating metal has
A higher reduction potential
B lower reduction potential
C same potential
D no potential
It oxidizes preferentially.
Cathodic protection prevents corrosion by
A oxidizing the metal
B making metal the cathode
C making metal the anode
D removing oxygen
Cathode does not undergo oxidation.
Fuel cells convert
A chemical energy directly into electrical energy
B heat into electricity
C electricity into chemical energy
D nuclear energy into electricity
Fuel cells are highly efficient electrochemical devices.
In hydrogen-oxygen fuel cell, the fuel is
A oxygen
B hydrogen
C water
D air
Hydrogen is oxidized at anode.
In a fuel cell, oxygen is reduced at
A anode
B cathode
C electrolyte
D wire
Oxygen gains electrons at cathode.
Corrosion is an example of
A electrolytic cell
B galvanic cell
C fuel cell
D dry cell
Corrosion involves spontaneous electrochemical reactions.
The flow of electrons during corrosion occurs from
A cathode to anode
B anode to cathode
C salt bridge to metal
D electrolyte to air
Same as any galvanic cell.
One method to prevent corrosion is
A increasing humidity
B painting the surface
C scratching the surface
D heating the metal
Paint acts as a barrier to air and moisture