When balancing redox reactions by the ion–electron method, the key idea is to split the overall reaction into oxidation and reduction half-reactions so that electron transfer can be balanced explicitly.
A Split into half-reactions
B Add catalyst first
C Remove spectators first
D Balance oxygen last
The ion–electron method becomes systematic by separating oxidation and reduction. Each half-reaction is balanced for atoms and charge, then combined by equalizing electrons, ensuring conservation of mass and charge.
In the oxidation number method, the total increase in oxidation number must be equal to the total decrease in oxidation number for a correctly balanced redox equation.
A Must be unequal
B Must be equal
C Must be doubled
D Must be ignored
Redox balancing is based on electron conservation. Increase in oxidation number represents electrons lost, and decrease represents electrons gained, so the net electrons exchanged must balance exactly.
In acidic medium, after balancing oxygen atoms by adding water, the hydrogen atoms are balanced by adding hydrogen ions on the appropriate side.
A Add OH⁻ ions
B Add H⁺ ions
C Add O₂ gas
D Add electrons only
Acidic medium contains excess H⁺, so it is chemically consistent to use H⁺ to balance hydrogen atoms after oxygen is balanced with water. Charge balancing is then completed using electrons.
In basic medium redox balancing, any H⁺ ions that appear can be removed by adding an equal number of OH⁻ ions to both sides, producing water.
A Add H⁺ to both sides
B Add OH⁻ to both sides
C Add salt bridge ions
D Add metal ions
Basic medium requires OH⁻ as the balancing species. Adding OH⁻ neutralizes H⁺ to form water, keeping the final balanced equation consistent with alkaline conditions.
The oxidation number of oxygen is generally −2 in most compounds, except in peroxides and superoxides where it becomes less negative.
A −2 always
B −1 in peroxides
C +2 in oxides
D 0 in compounds
Oxygen usually has oxidation number −2, but in peroxides (like H₂O₂) it is −1 because the O–O bond shares electrons differently. This exception is frequently tested.
The oxidation number of hydrogen is typically +1 when bonded to non-metals, but it becomes −1 in metal hydrides due to higher electropositivity of metals.
A +1 in hydrides
B −1 in metal hydrides
C 0 in acids
D −2 in water
In metal hydrides such as NaH, hydrogen behaves like a hydride ion (H⁻) because the metal donates electron density. Recognizing this helps correct oxidation state calculations in redox problems.
A disproportionation reaction is one in which the same element in one oxidation state is simultaneously oxidized and reduced, forming products with different oxidation states.
A Combination reaction
B Displacement reaction
C Disproportionation
D Neutralization
Disproportionation involves internal redox: the same species acts as both oxidizing and reducing agent. A classic idea is one oxidation state splitting into higher and lower states in products.
In a comproportionation reaction, two species of the same element in different oxidation states react to form a product where that element has an intermediate oxidation state.
A Decomposition
B Comproportionation
C Hydrolysis
D Polymerization
Comproportionation is the opposite of disproportionation. It combines different oxidation states of the same element to reach an intermediate state, showing electron transfer between forms.
In a galvanic cell, the salt bridge mainly allows ions to migrate so that each half-cell remains electrically neutral as oxidation and reduction proceed.
A Creates electrons
B Maintains neutrality
C Increases EMF
D Stops ion flow
Without a salt bridge, charge would build up: anode solution becomes positively charged and cathode solution becomes negatively charged. Ion migration prevents this and keeps the cell operating continuously.
A porous partition or salt bridge reduces liquid junction potential and helps prevent direct mixing of the electrolytes while still permitting ionic conduction.
A Blocks all ions
B Prevents electrolyte mixing
C Produces oxygen gas
D Converts ions to atoms
The salt bridge limits bulk mixing while allowing ions to move. This avoids unwanted side reactions and maintains a stable potential by minimizing junction potential differences.
In a Daniell cell, zinc undergoes oxidation at the anode and copper ions undergo reduction at the cathode, producing an overall spontaneous cell reaction.
A Zn reduces Cu²⁺
B Cu reduces Zn²⁺
C Zn oxidizes Cu
D Cu oxidizes Zn
Zinc loses electrons to form Zn²⁺ (oxidation) and those electrons reduce Cu²⁺ to Cu at the cathode. This electron transfer is the basis of the Daniell cell’s electrical output.
Standard cell potential for a galvanic cell under standard conditions is calculated using standard reduction potentials as E°cell = E°cathode − E°anode.
A E°anode − E°cathode
B E°cathode − E°anode
C E°cathode + E°anode
D E°anode + E°cathode
Reduction potentials are tabulated as reductions. The cathode is the site of reduction and anode is the site of oxidation, so subtracting anode potential from cathode gives the net driving potential.
If E°cell is positive, the corresponding redox reaction is spontaneous under standard conditions and can be used to generate electrical energy in a galvanic cell.
A Always non-spontaneous
B Spontaneous
C Always at equilibrium
D Impossible to predict
A positive E°cell implies ΔG° is negative because ΔG° = −nFE°cell. Negative free energy change indicates spontaneity, making the cell capable of producing electrical work.
The relationship between standard Gibbs free energy change and standard cell potential is given by ΔG° = −nF E°cell, linking thermodynamics with electrochemistry.
A ΔG° = +nFE°
B ΔG° = −nFE°
C ΔG° = nRT/E°
D ΔG° = F/nE°
This equation directly connects the electrical work obtainable from a cell with thermodynamic driving force. The negative sign shows that a positive cell potential corresponds to a negative free energy change.
The equilibrium constant for a redox reaction in an electrochemical cell is related to E°cell through the equation E° = (0.0591/n) log K at 298 K.
A Not related
B Directly related via log K
C Related to density
D Related to viscosity
Larger E°cell means larger K, indicating reaction proceeds strongly toward products. This relation helps connect electrochemical measurements with chemical equilibrium concepts.
According to the Nernst equation, increasing the concentration of products generally decreases the cell potential because the reaction quotient increases.
A Increases Ecell
B Decreases Ecell
C Makes Ecell zero always
D Has no effect
In Nernst equation, E = E° − (0.0591/n) log Q. If products increase, Q increases, log Q increases, and E decreases, reducing the driving force of the cell.
For a concentration cell, EMF becomes zero when the concentrations in both half-cells become equal because there is no driving force for ion movement.
A EMF becomes maximum
B EMF becomes negative
C EMF becomes zero
D EMF becomes infinite
Concentration cells derive EMF solely from concentration difference. When concentrations equalize, Q becomes 1, log Q becomes 0, and the potential difference disappears.
A reference electrode is used in electrochemistry because absolute electrode potentials cannot be measured directly and must be determined relative to a standard.
A Absolute values are easy
B Relative measurement required
C Potentials are constant always
D Electrodes do not react
Only potential differences are measurable. Therefore, standard electrodes like SHE are used to establish a reference scale so that other electrode potentials can be compared reliably.
Conductance (G) is the reciprocal of resistance (R), meaning a conductor with higher resistance will show lower conductance for the same conditions.
A G = R
B G = 1/R
C G = R²
D G = R/2
Conductance represents ease of current flow. Since resistance opposes current, conductance is defined as inverse of resistance, which is a basic relationship in electrical measurements.
The SI unit of conductance is siemens, which is equivalent to ohm⁻¹ and represents how easily current passes through a conductor or solution.
A Ohm
B Siemens
C Volt
D Farad
Conductance has unit S (siemens), equal to reciprocal of ohm. In electrochemistry, conductance helps quantify how well electrolytic solutions carry electric current.
Specific conductance of a solution typically increases with dilution for very concentrated solutions due to decreased viscosity and improved ionic mobility.
A Always decreases
B Always increases
C Often decreases on dilution
D Independent of ions
Specific conductance depends on number of ions per unit volume. Dilution reduces ion concentration per volume, so specific conductance generally decreases even though mobility may increase.
Molar conductance increases on dilution because it accounts for conductance contributed by all ions produced from one mole of electrolyte, which increases as dissociation improves.
A Decreases always
B Increases on dilution
C Becomes negative
D Becomes constant at all dilutions
On dilution, weak electrolytes ionize more and inter-ionic attractions decrease, allowing greater effective conductance per mole. Hence molar conductance rises, especially sharply for weak electrolytes.
For strong electrolytes, the increase in molar conductance with dilution is relatively small because they are already almost completely dissociated at higher concentrations.
A They dissociate more
B They are already dissociated
C They form precipitate
D They do not conduct
Strong electrolytes produce ions nearly fully even in concentrated solutions, so dilution mainly reduces ion–ion interactions, causing only a slight rise in molar conductance.
Kohlrausch’s law of independent ionic migration is most useful for calculating limiting molar conductance of weak electrolytes indirectly from strong electrolytes.
A Enthalpy calculation
B Limiting conductance calculation
C Density calculation
D pH calculation
Weak electrolytes cannot be fully ionized easily, so their limiting molar conductance is obtained by combining ionic contributions using Kohlrausch’s law from measurable strong electrolyte data.
The degree of dissociation of a weak electrolyte can be estimated using the ratio of its molar conductance at a given concentration to its limiting molar conductance.
A α = Λm/Λm°
B α = Λm°/Λm
C α = Λm × Λm°
D α = Λm − Λm°
For weak electrolytes, molar conductance increases with dissociation. The fraction ionized (α) is approximated by Λm divided by Λm°, helping quantify ionization at a given concentration.
The internal resistance of a cell arises due to resistance offered by electrolyte, electrodes, and other internal components, which reduces the terminal voltage under load.
A Increases terminal voltage
B Reduces terminal voltage
C Stops chemical reaction
D Makes EMF zero always
When current flows, voltage drop occurs inside the cell due to internal resistance. Thus terminal voltage becomes less than EMF, which is an important practical consideration in batteries.
A primary cell is designed for one-time use because its chemical reaction is not practically reversible under normal operating conditions.
A Fully reversible
B Not practically reversible
C Needs recharging
D Works only in acid
Primary cells like dry cells undergo irreversible reactions during discharge. They cannot be recharged efficiently, unlike secondary cells where reactions are reversible.
In a lead-acid battery during discharge, both electrodes get converted to lead sulfate while sulfuric acid concentration decreases.
A Acid concentration increases
B Acid concentration decreases
C No change in acid
D Acid disappears completely
During discharge, H₂SO₄ is consumed to form PbSO₄ at both electrodes, so electrolyte becomes more dilute. This change is used to estimate battery state using density measurements.
Lithium-ion batteries operate based on intercalation and de-intercalation of lithium ions between electrode materials during charging and discharging.
A Gas evolution
B Metal deposition
C Ion intercalation
D Acid neutralization
Lithium ions move between anode and cathode structures without major chemical breakdown of electrodes, enabling high efficiency and rechargeability, which is why these batteries dominate electronics.
Corrosion of iron is accelerated in the presence of electrolytes like dissolved salts because they increase conductivity and facilitate electrochemical reactions on the metal surface.
A Conductivity decreases
B Conductivity increases
C Oxygen is removed
D Reaction stops
Dissolved salts enhance ionic conduction in water, allowing faster electron and ion movement in corrosion cells, thereby increasing corrosion rate under practical environmental conditions.
In electrochemical corrosion, the anodic region is the part of the metal surface where oxidation occurs and metal atoms convert into metal ions.
A Reduction occurs
B Oxidation occurs
C Neutralization occurs
D Precipitation occurs
At the anodic site, metal loses electrons (oxidation), forming metal ions. This is the actual dissolution step responsible for loss of metal during corrosion.
The cathodic reaction in rusting of iron commonly involves reduction of dissolved oxygen in the presence of water to produce hydroxide ions.
A Oxygen oxidation
B Oxygen reduction
C Iron reduction
D Water oxidation
At cathodic sites, oxygen gains electrons and forms OH⁻ in neutral or basic environments. These OH⁻ ions then react with Fe²⁺ to form hydroxides leading to rust formation.
Galvanization protects iron from corrosion by coating it with zinc, which acts as a sacrificial metal and oxidizes preferentially if the coating is damaged.
A Zinc reduces faster corrosion
B Zinc sacrificial protection
C Iron becomes noble
D Oxygen is removed completely
Zinc is more reactive than iron, so it oxidizes first and supplies electrons to iron, preventing iron oxidation. This works even if scratches expose the iron beneath.
Electroplating is a practical application of electrolysis where a thin layer of one metal is deposited on another metal surface to improve appearance or corrosion resistance.
A Mechanical coating
B Electrolytic deposition
C Thermal spraying
D Chemical reduction only
Electroplating uses an electrolytic cell so that metal ions in solution are reduced at the cathode and deposited as a coherent metal layer, widely used in industry.
In electrolysis, the cathode is the electrode where cations are discharged by gaining electrons, leading to deposition or gas evolution depending on conditions.
A Cations lose electrons
B Cations gain electrons
C Anions gain electrons
D Anions lose protons
Cathode is the site of reduction. Positively charged ions move toward the cathode and accept electrons, producing neutral atoms or molecules such as hydrogen gas.
The amount of substance deposited during electrolysis depends on the total charge passed, which itself depends on current and time according to Q = It.
A Q = I/t
B Q = It
C Q = I²t
D Q = t/I
Charge is the product of current and time. Since deposition is proportional to charge (Faraday’s law), controlling current and duration controls the deposited mass quantitatively.
If a redox reaction involves transfer of 2 moles of electrons, then the value of n used in ΔG = −nFE is 2 for that reaction.
A n = 1
B n = 2
C n = 4
D n = 6
In electrochemistry, n represents moles of electrons transferred per mole of overall reaction. Correct determination of n is essential for accurate ΔG and equilibrium calculations.
In a galvanic cell, the anode is negative and cathode is positive because electrons are produced at the anode and consumed at the cathode.
A Anode positive
B Cathode negative
C Anode negative
D Both positive
In galvanic cells, oxidation at anode releases electrons, making it electron-rich and negative. Cathode attracts electrons for reduction, making it positive in the external circuit.
In an electrolytic cell, the polarity reverses compared to galvanic cells because an external power source forces electrons toward the cathode.
A Anode is negative
B Cathode is negative
C Both are neutral
D Salt bridge sets polarity
In electrolysis, the cathode is connected to the negative terminal of the power supply, so it receives electrons and becomes the site of reduction, while the anode becomes positive.
The Nernst equation shows that for reactions where n is larger, the potential change with concentration becomes smaller because the concentration term is divided by n.
A Change becomes larger
B Change becomes smaller
C Change becomes random
D Change becomes infinite
The Nernst concentration correction is (0.0591/n) log Q at 298 K. A larger n reduces the magnitude of this term, making potential less sensitive to concentration changes.
If the reaction quotient Q equals 1, then log Q becomes zero and the cell potential equals the standard cell potential at that temperature.
A E = 0
B E = E°
C E becomes negative always
D E becomes infinite
When Q = 1, the system corresponds to standard-like ratio of species in the Nernst expression, so the logarithmic correction becomes zero and E equals E°.
A higher oxidation potential for a given half-cell implies a greater tendency for that electrode to undergo oxidation compared to another electrode.
A Less oxidation tendency
B Greater oxidation tendency
C No relation to oxidation
D Only depends on mass
Oxidation potential is essentially the negative of reduction potential for the same half-reaction. A higher oxidation potential indicates a stronger tendency to lose electrons and be oxidized.
In conductivity measurements, the cell constant depends on the geometry of the conductivity cell, typically expressed as distance between electrodes divided by electrode area.
A Depends only on solution
B Depends on geometry
C Depends only on temperature
D Depends on battery type
Cell constant corrects measured conductance to specific conductance by accounting for electrode spacing and area. It must be determined by calibration with standard solutions.
Specific resistance of a solution is the reciprocal of specific conductance, showing that better conducting solutions have lower resistivity.
A Directly proportional
B Reciprocal relation
C Unrelated concepts
D Equal always
Specific conductance (κ) measures conducting ability, while resistivity (ρ) measures opposition. They are reciprocals: ρ = 1/κ, a key relation in conductance numericals.
When a battery delivers current for a long time at nearly constant voltage, it is preferred for practical use because it ensures stable performance of devices.
A Unstable voltage preferred
B Constant voltage preferred
C High corrosion preferred
D Low capacity preferred
Devices require stable supply voltage for proper operation. Batteries designed for flat discharge curves maintain nearly constant potential, improving efficiency and reducing performance fluctuations.
Corrosion can be reduced by applying protective coatings like paint because such coatings physically block oxygen and moisture from reaching the metal surface.
A Increase metal exposure
B Block oxygen and moisture
C Produce more ions
D Increase conductivity
Paint and similar coatings act as barriers, preventing contact between metal and corrosive environment. Without water and oxygen access, electrochemical corrosion reactions slow significantly.
In corrosion, formation of a small anodic area with a large cathodic area generally increases corrosion rate because current density at the anodic site becomes high.
A Corrosion decreases
B Corrosion increases
C No change occurs
D Corrosion stops fully
Corrosion rate depends on anodic current density. A small anode must supply electrons for a large cathode, concentrating dissolution at the anode and accelerating localized corrosion.
A battery with higher capacity is one that can supply a given current for a longer time, indicating larger stored charge available for discharge.
A Lower stored charge
B Larger stored charge
C Lower efficiency always
D Higher resistance always
Capacity relates to total charge deliverable (ampere-hour). More capacity means the battery can run devices longer at the same current, reflecting greater usable electrochemical energy.
The term electromotive force refers to the maximum potential difference between electrodes of a cell when no current is flowing, representing the ideal driving force.
A Voltage under load
B Maximum no-current voltage
C Current through wire
D Resistance of electrolyte
EMF is measured under open-circuit conditions. When current flows, internal resistance causes voltage drop, so terminal voltage becomes less than EMF.
An effective corrosion prevention strategy in underground pipelines is cathodic protection using an external power source, which continuously supplies electrons to keep the pipeline cathodic.
A Increase anodic dissolution
B Cathodic protection
C Use only water coating
D Increase oxygen exposure
Impressed current cathodic protection forces the structure to behave as a cathode, preventing metal oxidation. Continuous electron supply counters corrosion reactions, making it highly effective in pipelines.