The rate of a chemical reaction is defined as the change in
A temperature per unit time
B concentration per unit time
C pressure per unit volume
D energy per unit mass
Rate measures how fast reactant concentration decreases or product concentration increases with time.
The SI unit of rate of reaction is
A mol L⁻¹
B mol s⁻¹
C mol L⁻¹ s⁻¹
D L mol⁻¹ s⁻¹
Rate = change in concentration (mol L⁻¹) per time (s).
The rate law of a reaction is determined by
A balanced chemical equation
B experimental observation
C enthalpy change
D equilibrium constant
Rate law cannot be predicted reliably from equation; it is found experimentally.
Which factor does not directly determine reaction rate
A Temperature
B Catalyst
C Concentration
D Standard enthalpy change (ΔH°)
ΔH is thermodynamic; rate is kinetic and depends on pathway and activation energy.
The order of reaction equals
A sum of stoichiometric coefficients
B sum of powers of concentration terms in rate law
C number of products
D number of steps in mechanism
Order = exponents of concentration terms in rate expression.
For Rate = k[A]²[B], the order is
A 2
B 3
C 1
D 4
Total order = 2 + 1 = 3.
Molecularity is defined for
A overall reaction only
B elementary step only
C reversible reaction only
D equilibrium state
Molecularity counts molecules colliding in a single elementary step.
Molecularity can never be
A 1
B 2
C 3
D fractional
Molecularity is always a whole number (1,2,3…).
Order of reaction can be fractional in
A elementary reactions
B complex reactions
C unimolecular reactions only
D only zero-order reactions
Fractional order arises from multi-step mechanisms/steady-state behavior.
For a zero-order reaction, rate is independent of
A temperature
B catalyst
C reactant concentration
D nature of reactant
Zero-order rate = k (constant), not dependent on [A].
Unit of k for a zero-order reaction is
A s⁻¹
B mol L⁻¹ s⁻¹
C L mol⁻¹ s⁻¹
D L² mol⁻² s⁻¹
For zero order, rate = k, so k has same unit as rate.
Unit of k for a first-order reaction is
A mol L⁻¹ s⁻¹
B s⁻¹
C L mol⁻¹ s⁻¹
D L² mol⁻² s⁻¹
Rate = k[A] ⇒ k = rate/[A] ⇒ (mol L⁻¹ s⁻¹)/(mol L⁻¹) = s⁻¹.
Unit of k for a second-order reaction (Rate = k[A]²) is
A s⁻¹
B mol L⁻¹ s⁻¹
C L mol⁻¹ s⁻¹
D L² mol⁻² s⁻¹
k = rate/[A]² ⇒ (mol L⁻¹ s⁻¹)/(mol² L⁻²) = L mol⁻¹ s⁻¹.
Which graph is linear for a zero-order reaction
A [A] vs t
B log[A] vs t
C 1/[A] vs t
D ln(1/[A]) vs t
Zero order integrated form: [A] = [A]₀ − kt (straight line).
Which graph is linear for a first-order reaction
A [A] vs t
B log[A] vs t
C [A]² vs t
D 1/[A] vs t
First order integrated form: log[A] = log[A]₀ − (k/2.303)t.
Which graph is linear for a second-order reaction (Rate = k[A]²)
A [A] vs t
B log[A] vs t
C 1/[A] vs t
D 1/[A]² vs t
Second order integrated form: 1/[A] = 1/[A]₀ + kt.
Order of reaction is
A always equal to molecularity
B always equal to stoichiometric coefficients
C determined experimentally
D always 1 or 2 only
Order depends on mechanism; measured from rate law.
Molecularity of a bimolecular elementary step is
A 1
B 2
C 3
D 0
Two species collide in the elementary step.
A termolecular step means collision of
A 1 molecule
B 2 molecules
C 3 molecules
D many molecules
Termolecular means three reactant species involved in the same step (rare).
Rate constant k depends on
A concentration only
B temperature and nature of reactants
C initial amount of reactant only
D stoichiometric coefficients only
k changes with temperature (Arrhenius) and reaction pathway/nature.
In general, for reactions of gases, increasing pressure at constant T usually
A decreases rate always
B increases rate by increasing concentration
C has no effect ever
D changes ΔH only
Higher pressure increases effective concentration/collision frequency.
If doubling [A] doubles rate, order w.r.t A is
A 0
B 1
C 2
D 3
Rate ∝ [A]¹ if doubling [A] doubles rate.
If doubling [A] makes rate four times, order w.r.t A is
A 0
B 1
C 2
D 1/2
Rate ∝ [A]² gives factor 2² = 4.
If doubling [A] makes rate unchanged, order w.r.t A is
A 0
B 1
C 2
D −1
Zero order: rate independent of [A].
A negative order indicates that increasing concentration
A increases rate
B decreases rate
C has no effect
D increases equilibrium constant
Negative order can occur in complex mechanisms (e.g., inhibition by reactant).
For Rate = k[A]^0[B]^1, overall order is
A 0
B 1
C 2
D 3
Total order = 0 + 1 = 1.
Which statement is correct
A Molecularity may be zero
B Order is always integer
C Molecularity is always integer
D Order is always same as molecularity
Molecularity counts molecules in one step; must be whole number.
If a reaction is written as A + B → products but rate = k[A], then
A B is catalyst
B B is spectator or in excess or involved in fast equilibrium
C order must be 2
D molecularity must be 2 always
Rate law indicates effective dependence; mechanism/excess can hide B dependence.
For an elementary reaction A → products, molecularity is
A 0
B 1
C 2
D 3
One molecule decomposes.
For an elementary reaction A + A → products, molecularity is
A 1
B 2
C 3
D cannot be defined
Two molecules collide (even if same species).
Rate law is valid for
A all temperatures without change
B a given mechanism and conditions
C only at equilibrium
D only for elementary reactions
Change in conditions/mechanism may change observed rate law.
A catalyst changes the rate primarily by
A changing ΔG°
B changing equilibrium constant
C lowering activation energy
D increasing reactant concentration permanently
Catalyst provides alternate pathway with lower Ea.
Reaction order is obtained from
A integrated rate equation slope experiments
B balanced equation directly
C Hess law
D heat of reaction
Commonly from initial rates or linear plots of integrated forms.
If a plot of rate vs [A] is a straight line passing through origin, order is likely
A zero
B first
C second
D fractional
rate ∝ [A] indicates first order.
If a plot of rate vs [A] is horizontal line, order is
A zero
B first
C second
D third
Rate constant regardless of concentration indicates zero order.
In general, increasing surface area of a solid reactant increases rate because
A it lowers ΔH
B it increases number of collision sites
C it decreases temperature
D it changes order always
More exposed surface → more effective collisions.
The rate law for elementary step NO + O₃ → NO₂ + O₂ is expected as
A k[NO]
B k[O₃]
C k[NO][O₃]
D k[NO]²[O₃]
Elementary bimolecular step ⇒ first order in each reactant.
A reaction shows rate = k, it means the reaction is
A first order
B second order
C zero order
D third order
Rate independent of concentration implies zero order.
The slowest step in a mechanism is called
A initiation step
B propagation step
C rate-determining step
D termination step
It controls overall reaction rate.
For a multi-step reaction, overall rate law is mainly controlled by
A fastest step
B average of all steps
C rate-determining step and prior equilibria
D product stability only
Slow step + pre-equilibrium commonly determines rate expression.
If order is 1.5 overall, it indicates
A reaction is elementary
B reaction is complex
C reaction is zero order
D molecularity is 1.5
Fractional order comes from complex mechanism.
If Rate = k[A]⁻¹, then increasing [A] will
A increase rate
B decrease rate
C not change rate
D increase k
Negative order means inverse dependence.
The dimensional formula of rate constant for nth order reaction is
A (concentration)⁻ⁿ time⁻¹
B (concentration)^(1−n) time⁻¹
C (concentration)^(n−1) time
D time only
Rate = k[A]^n ⇒ k = rate/[A]^n ⇒ (conc/time)/(conc^n) = conc^(1−n)/time.
A reaction is first order overall if
A rate doubles when both concentrations double
B rate increases 4 times when concentration doubles
C rate does not change when concentration changes
D rate increases 8 times when concentration doubles
Overall order 1 implies rate ∝ concentration¹.
In kinetics, “initial rate method” uses
A equilibrium concentrations
B rate at very long time
C rate at very small time near start
D average rate only
Initial rates avoid complications like reverse reaction and product effects.
The order with respect to a reactant is the power of its concentration in
A equilibrium expression
B rate expression
C pH expression
D enthalpy equation
It is defined from rate law.
If reaction is 2nd order overall, then doubling all reactant concentrations makes rate
A 2 times
B 4 times
C 8 times
D unchanged
Rate ∝ (concentration)²; factor 2² = 4.
When a reactant is in very large excess, a higher-order reaction may appear as
A zero order
B pseudo-first order
C third order always
D not measurable
Excess reactant remains nearly constant; rate seems first order in limiting reactant.
For pseudo-first order, the observed constant is
A true k only
B k’ = k[B]ⁿ (B constant)
C independent of temperature
D same as equilibrium constant
Constant excess concentration gets absorbed into a new constant k’.
Which is correct about order
A cannot be zero
B can be zero, fractional, or even negative
C always equal to molecularity
D always obtained from balanced equation
Order comes from experimental kinetics and can take these values in complex reactions.