Arrhenius equation is
A k = [A]₀ − kt
B k = Ae^(−Ea/RT)
C 1/[A] = 1/[A]₀ + kt
D log[A] = log[A]₀ − (k/2.303)t
It relates rate constant to temperature and activation energy.
In Arrhenius equation, A is called
A activation energy
B frequency (pre-exponential) factor
C gas constant
D rate of reaction
A represents collision frequency and proper orientation probability.
Activation energy (Ea) is
A energy released in reaction
B energy needed to form activated complex
C total energy of products
D total energy of reactants only
It is minimum energy barrier for effective collisions.
Rate constant increases with temperature mainly because
A equilibrium constant increases always
B more molecules cross activation barrier
C ΔH becomes more negative
D reaction becomes zero order
Higher T increases fraction with energy ≥ Ea.
If Ea is high, reaction is generally
A very fast at low temperature
B slow at low temperature
C independent of temperature
D always spontaneous
Few molecules have enough energy when Ea is large.
A catalyst increases rate by
A increasing Ea
B lowering Ea
C increasing ΔG°
D changing Kc
Alternate pathway with lower barrier.
Catalyst does not change
A rate
B activation energy
C equilibrium constant
D reaction pathway
Catalyst affects kinetics, not thermodynamic equilibrium.
The Arrhenius plot is a graph of
A k vs T
B log k vs 1/T
C k vs 1/T
D log k vs T
It gives straight line; slope related to Ea.
For Arrhenius plot, slope equals
A +Ea/2.303R
B −Ea/2.303R
C −2.303R/Ea
D +2.303R/Ea
log k = log A − (Ea/2.303R)(1/T).
If slope becomes more negative, Ea is
A smaller
B larger
C zero
D unrelated
Slope magnitude ∝ Ea.
Increasing temperature by 10°C often increases rate by about
A 1.1 to 1.2 times
B 2 to 3 times
C 10 times
D no change
Typical empirical temperature coefficient (varies).
Collision theory states that reaction occurs when molecules
A collide with any energy
B collide with sufficient energy and proper orientation
C are at equilibrium
D have low energy only
Effective collisions require energy ≥ Ea and correct orientation.
In collision theory, the fraction of effective collisions increases with
A decreasing temperature
B increasing temperature
C decreasing concentration
D decreasing surface area
More molecules exceed Ea.
The energy profile diagram shows activated complex at
A lowest point
B highest point on curve
C at reactant level always
D at product level always
Transition state is peak energy point.
Ea for catalysed reaction is
A greater than uncatalysed
B equal to uncatalysed
C less than uncatalysed
D independent of catalyst
Catalyst lowers activation energy.
If temperature increases, rate constant k
A decreases always
B increases generally
C becomes zero
D becomes negative
Arrhenius dependence.
The equation log(k2/k1) = (Ea/2.303R)(1/T1 − 1/T2) is used to find
A order
B molecularity
C Ea using two temperatures
D concentration at time t
Two-point Arrhenius form.
If k doubles when T rises slightly, it implies
A Ea is zero
B Ea is significant and temperature sensitive
C reaction is zero order
D equilibrium constant doubled
Strong T-dependence indicates meaningful Ea.
A large A (frequency factor) suggests
A very low collision frequency
B very high collision frequency/orientation probability
C low temperature only
D zero activation energy
A reflects frequency and orientation contributions.
If Ea = 0, then k is
A zero
B independent of temperature (approximately)
C negative
D equal to R
k = Ae^(0) = A, so minimal T effect.
In Arrhenius equation, if temperature increases, the value of e^(−Ea/RT) generally
A decreases
B increases
C remains constant
D becomes negative
As T increases, Ea/RT decreases, so the negative exponent becomes less negative, increasing the exponential term and thus k.
For two temperatures T1 and T2 (T2 > T1), the ratio k2/k1 is
A always < 1
B always > 1
C always = 1
D unpredictable for any reaction
For most reactions, k increases with temperature, so k2 > k1.
If Ea is very large, the reaction rate becomes
A almost independent of temperature
B highly sensitive to temperature
C independent of concentration
D always zero order
Larger Ea means a small increase in T causes a large increase in molecules crossing the barrier.
For Arrhenius plot log k vs 1/T, the intercept equals
A log Ea
B log A
C −Ea/R
D A/Ea
log k = log A − (Ea/2.303R)(1/T). When 1/T → 0, log k → log A.
A catalyst increases reaction rate mainly by increasing
A ΔH of reaction
B Ea of reaction
C number of effective collisions
D equilibrium constant
Catalyst lowers Ea, increasing the fraction of collisions with energy ≥ Ea (effective collisions).
For the same reaction at the same temperature, the catalysed reaction has
A higher Ea and higher k
B lower Ea and higher k
C lower Ea and lower k
D higher Ea and lower k
Lower activation energy gives larger rate constant at same temperature.
The “activated complex” is
A most stable species
B intermediate with minimum energy
C unstable high-energy transition state
D always a product
It is formed at the peak of energy barrier and quickly converts to products or reactants.
The collision frequency increases mainly with
A decreasing temperature
B increasing temperature
C decreasing volume (in gases)
D both B and C
Higher T increases molecular speed; lower volume increases concentration—both increase collision frequency.
According to collision theory, reaction rate depends on
A only collision frequency
B only activation energy
C collision frequency and fraction of effective collisions
D only entropy change
Not every collision leads to reaction; only effective collisions matter.
For a reaction with very small Ea, increasing temperature will cause
A very large increase in rate
B very small increase in rate
C rate to decrease
D rate to become zero
When Ea is small, fraction of molecules above Ea is already high; temperature change has less effect.
If a reaction has k = Ae^(−Ea/RT), then at very high T, k approaches
A 0
B A
C −A
D Ea
As T becomes very large, Ea/RT → 0, so e^0 = 1 ⇒ k ≈ A.
If the slope of Arrhenius plot is −5000 K, then Ea is (R = 8.314 J mol⁻¹ K⁻¹)
A 9.6 kJ mol⁻¹
B 20.7 kJ mol⁻¹
C 95.7 kJ mol⁻¹
D 5.0 kJ mol⁻¹
slope = −Ea/(2.303R).
Ea = 2.303R × 5000 = 2.303×8.314×5000 ≈ 95700 J mol⁻¹ = 95.7 kJ mol⁻¹.
The temperature coefficient generally represents the factor by which rate increases for
A 1°C rise
B 5°C rise
C 10°C rise
D 100°C rise
Temperature coefficient is commonly defined for a 10°C rise.
The effect of temperature on reaction rate is stronger for reactions with
A lower Ea
B higher Ea
C zero order only
D equilibrium reactions only
High Ea reactions show larger change in k with temperature.
Which one is NOT a kinetic factor
A Activation energy
B Reaction mechanism
C Rate constant
D Standard Gibbs free energy change (ΔG°)
ΔG° is thermodynamic (feasibility), not speed.
Increasing temperature increases reaction rate mainly because
A reactants become heavier
B number of collisions decreases
C more molecules have energy ≥ Ea
D equilibrium constant always increases
Maxwell–Boltzmann distribution shifts, increasing effective collisions.
If k1 and k2 are rate constants at T1 and T2, then log(k2/k1) equals
A (Ea/2.303R)(1/T2 − 1/T1)
B (Ea/2.303R)(1/T1 − 1/T2)
C (2.303R/Ea)(T2 − T1)
D (Ea/R)(T2/T1)
Standard two-temperature Arrhenius form.
A catalyst provides
A same pathway with higher Ea
B alternate pathway with lower Ea
C alternate pathway with higher ΔH
D alternate pathway changing Kc
Rate increases because barrier decreases; equilibrium unchanged.
For an endothermic reaction, increasing temperature generally
A decreases k
B increases k
C makes k zero
D makes k negative
Arrhenius relation holds regardless of endo/exothermic nature.
If Ea is expressed in kJ mol⁻¹, then R should be used as
A 8.314 kJ mol⁻¹ K⁻¹
B 8.314 J mol⁻¹ K⁻¹ with conversion
C 0.0821 L atm mol⁻¹ K⁻¹ only
D 1.987 cal mol⁻¹ only
Keep units consistent (convert Ea to J mol⁻¹ or R to kJ mol⁻¹ K⁻¹).
The energy distribution of molecules in a gas is described by
A Boyle’s law
B Maxwell–Boltzmann distribution
C Raoult’s law
D Faraday law
It shows fraction of molecules at various energies and explains temperature effect on rate.
On increasing temperature, the Maxwell–Boltzmann curve
A becomes taller and narrower
B shifts to higher energy and flattens
C shifts to lower energy and sharpens
D remains unchanged
Higher T spreads energies and increases high-energy fraction.
The orientation factor in collision theory is included in
A Ea
B A (pre-exponential factor)
C R
D T
A includes collision frequency and proper orientation probability.
If a catalyst doubles the rate constant at same T, it implies
A Ea is unchanged
B Ea is reduced (generally)
C ΔH becomes zero
D Kc becomes double
Increased k at same T commonly results from reduced Ea.
A reaction has Arrhenius plot slope = −4000 K. Ea is approximately
A 76.7 kJ mol⁻¹
B 36.7 kJ mol⁻¹
C 7.67 kJ mol⁻¹
D 4000 kJ mol⁻¹
Ea = 2.303R×4000 ≈ 2.303×8.314×4000 ≈ 76600 J = 76.6 kJ mol⁻¹.
If k increases from 0.10 s⁻¹ to 0.20 s⁻¹, then rate at same concentration becomes
A half
B double
C four times
D unchanged
For same rate law and concentration, rate ∝ k.
In an energy profile diagram, catalyst effect is shown by
A raising product energy
B lowering peak height
C increasing ΔH
D shifting reactant level upward
Peak height corresponds to activation energy.
Ea for forward and backward reactions differ because
A ΔH changes sign
B equilibrium constant changes
C reactant concentration changes
D pressure changes
Ea(reverse) = Ea(forward) − ΔH (sign convention matters).
A reaction rate becomes extremely high at high temperature mainly because
A ΔG becomes positive
B fraction of molecules above Ea increases rapidly
C molecularity changes
D order becomes zero
Exponential Arrhenius dependence can produce huge rate rise.
The correct statement is
A catalyst increases Kc
B catalyst lowers Ea for both forward and backward equally
C catalyst changes ΔH
D catalyst is consumed permanently
Catalyst provides alternate path for both directions; equilibrium unchanged.