Class 12 Chemistry · NCERT chapter 4 · 7 topics · free to read
Rate of Reaction
Rate of reaction measures how fast reactants are consumed or products are formed. Average rate = Δ[concentration]/Δt over a time interval. Instantaneous rate = −d[R]/dt = d[P]/dt at a specific instant (slope of concentration-time curve). Rate depends on concentration, temperature, catalyst, and surface area.
Rate = - 1 a d[A] dt = - 1 b d[B] dt = 1 c d[C] dt = 1 d d[D] dt
Average Rate
The average rate of reaction is the change in concentration of a reactant or product divided by the time interval. It gives a rough measure of 'how fast' over a finite period — like average speed on a road trip.
For a reaction A arrow B , the average rate over time interval t = t_2 - t_1 is:
Average rate = - [A] t = - [A]_ t_2 - [A]_ t_1 t_2 - t_1 = [B] t
Negative sign for reactants because [A] decreases — the minus sign makes rate positive
No negative sign for products because [B] increases naturally
Units: mol L ^ -1 s ^ -1 (or mol L ^ -1 min ^ -1 , depending on the time unit)
Average rate decreases as reaction progresses — reactants get used up, collisions become less frequent
Instantaneous Rate
Instantaneous rate is the rate at a specific instant — the slope of the tangent to the concentration-time curve at that point. As Δt → 0, the average rate becomes the instantaneous rate.
The instantaneous rate is the rate at one precise moment in time. Mathematically, it's the limit of the average rate as t 0 :
Instantaneous rate = - d[A] dt = _ t 0 (- [A] t )
Graphically, draw a tangent to the [A] vs t curve at the desired time. The slope of that tangent (with a sign correction for reactants) is the instantaneous rate.
Instantaneous rate at t = 0 is called the initial rate — the fastest rate (maximum reactant concentration)
Instantaneous rate is always measured from the slope of the tangent to the concentration-time graph
NCERT specifies: when we simply say 'rate', we usually mean instantaneous rate
Initial rate method is used experimentally to determine order of reaction
Rate & Stoichiometry
For a general reaction aA + bB → cC + dD, the rates of disappearance of reactants and appearance of products are related by stoichiometric coefficients. Dividing by the coefficient gives a unique rate for the overall reaction.
For the reaction aA + bB arrow cC + dD , species disappear and appear at different absolute rates depending on their stoichiometric coefficients. To define a single unique rate, we divide each rate by its coefficient:
Rate = - 1 a d[A] dt = - 1 b d[B] dt = + 1 c d[C] dt = + 1 d d[D] dt
Five factors affect reaction rate: concentration (more collisions), temperature (faster molecules), catalyst (lower activation energy), surface area (more contact), and nature of reactants.
The rate of a reaction can be changed by tweaking five factors. Each one works by affecting either the frequency of collisions or the fraction of molecules with enough energy to react.
Concentration — Higher concentration → more molecules per unit volume → more frequent collisions → faster rate
Temperature — Higher temperature → molecules move faster → more energetic collisions → exponentially faster rate
Catalyst — Provides an alternative pathway with lower activation energy → more molecules can cross the barrier → faster rate without being consumed
Surface area — Grinding a solid into powder exposes more surface → more contact with other reactants → faster rate
Nature of reactants — Ionic reactions are fast (no bonds to break), covalent reactions involving bond rearrangement are slower
Rate Law & Order
The rate law expresses rate as a function of concentrations: Rate = k[A]^n[B]^m. The exponents n, m are the order with respect to each reactant (determined experimentally, NOT from stoichiometry). Overall order = n + m. Molecularity is the number of molecules in an elementary step. Pseudo-first order reactions have one reactant in large excess.
Rate = k[A]^n[B]^m (overall order = n + m)
Rate Law
The rate law (or rate equation) relates the rate to concentrations of reactants raised to experimentally determined powers. The proportionality constant k is the rate constant — it absorbs temperature and catalyst effects.
For a reaction aA + bB arrow Products , the rate law is:
Rate = k[A]^n[B]^m
Here k is the rate constant (specific rate constant), and n , m are the orders with respect to A and B . The crucial point: n and m are found experimentally — they may or may not equal the stoichiometric coefficients a and b .
Rate law can only be determined experimentally — never from the balanced equation alone
Rate constant k depends on temperature and catalyst but NOT on concentration
Units of k depend on the overall order: for n th order, k has units mol ^ 1-n L ^ n-1 s ^ -1
Rate law is valid only for elementary reactions if written directly from stoichiometry
Order of Reaction
Order with respect to a reactant is its exponent in the rate law. Overall order is the sum of all exponents. Order can be 0, 1, 2, 3, or even fractional. It is always determined experimentally.
For Rate = k[A]^n[B]^m : - Order with respect to A = n - Order with respect to B = m - Overall order = n + m
Zero order ( n = 0 ): Rate is independent of [A] — doubling [A] has no effect on rate
First order ( n = 1 ): Rate is directly proportional to [A] — double [A] , double the rate
Second order ( n = 2 ): Rate ∝ [A]^2 — double [A] , rate quadruples
Third order ( n = 3 ): Rate ∝ [A]^3 — double [A] , rate increases 8×
Molecularity is the number of reacting species (atoms, ions, molecules) that collide simultaneously in an elementary reaction step. Unlike order, molecularity is theoretical, always a positive integer, and defined only for elementary steps.
Molecularity is the number of reacting species that must simultaneously collide for an elementary reaction to occur. It applies ONLY to elementary (single-step) reactions.
Bimolecular (molecularity = 2): Two species collide. Example: CH_3COOC_2H_5 + NaOH arrow products
Trimolecular (molecularity = 3): Three species collide simultaneously. Very rare — probability of three-body collision is extremely low
Molecularity > 3 is practically impossible — four molecules meeting at the exact same point with the right energy and orientation is near-zero probability
Experimental (measured)
Theoretical (counted from equation)
Can be 0, fractional, or negative
Always a positive integer (1, 2, or 3)
Defined for overall AND elementary reactions
Defined ONLY for elementary reactions
Can change with conditions
Fixed for a given elementary step
Sum of exponents in rate law
Number of colliding species
Pseudo-First Order
A pseudo-first order reaction is actually second (or higher) order, but one reactant is in such large excess that its concentration barely changes. Its concentration gets absorbed into k, making the reaction appear first order experimentally.
Consider: Rate = k[A][B] (second order overall). If [B] is in huge excess (say, water as solvent), then [B] barely changes during the reaction. We can treat [B] as a constant and absorb it into k :
Rate = k[A][B] = (k[B])_ const [A] = k'[A] where k' = k[B]
Now the rate law looks first order in A only. The true second-order rate constant k is hiding behind the pseudo-first-order constant k' = k[B] .
Inversion of cane sugar is the other classic example: C_ 12 H_ 22 O_ 11 + H_2O H^+ C_6H_ 12 O_6 + C_6H_ 12 O_6
The reactant in excess must be in large excess (10× or more) for the pseudo-order approximation to hold
Pseudo-first order constant k' has first-order units (s ^ -1 ), not second-order units
NCERT specifically mentions: pseudo-first order reactions are bimolecular but experimentally first order
Integrated Rate Equations
Integrated rate equations relate concentration to time directly (not rate to concentration). Zero order: [A] = [A₀] − kt (linear decrease). First order: ln([A₀]/[A]) = kt or [A] = [A₀]e^(−kt) (exponential decay). Second order: 1/[A] = 1/[A₀] + kt. Each gives a characteristic straight-line plot for determining order experimentally.
[A]_0 [A] = kt (first order)
Zero Order
For a zero-order reaction, rate is constant and independent of concentration. Integrating Rate = k gives [A] = [A₀] − kt. The plot of [A] vs t is a straight line with slope = −k.
For a zero-order reaction: Rate = k (constant, independent of [A] ). The rate doesn't care how much reactant is left — it keeps consuming at the same constant speed.
[A] = [A]_0 - kt
k = [A]_0 - [A] t Units: mol L^ -1 s^ -1
Plot of [A] vs t → straight line, slope = -k , intercept = [A]_0
Plot of Rate vs [A] → horizontal line (rate is constant)
Reaction completes when [A] = 0 , i.e., at t = [A]_0/k
Example: decomposition of NH _3 on hot platinum surface, photochemical reactions where light intensity controls rate
Units of k : mol L ^ -1 time ^ -1 (same as rate — because rate = k here)
First Order
For first-order reactions, integration gives ln([A₀]/[A]) = kt. The concentration decays exponentially. Plot of ln[A] vs t gives a straight line with slope −k. This is the most tested rate equation in NEET.
For a first-order reaction: Rate = k[A] . Separating variables and integrating:
[A]_0 [A] = kt [A] = [A]_0 - kt
k = 2.303 t [A]_0 [A] (log_ 10 form — most used in NEET)
[A] = [A]_0 e^ -kt (exponential decay form)
Plot of [A] vs t → straight line, slope = -k , intercept = [A]_0
Plot of [A] vs t → straight line, slope = -k/2.303
Units of k : s ^ -1 (or min ^ -1 , time ^ -1 ) — note: no concentration unit
First-order reactions never reach [A] = 0 — concentration asymptotically approaches zero
The form ( k = 2.303 t [A]_0 [A] ) is preferred in NEET because calculators use _ 10
Quick recall
A first-order reaction has k = 0.0693 min ^ -1 . How long for 75% completion?
75% done → 25% remains → [A] = 0.25[A]_0 . Using t = 2.303 k [A]_0 [A] = 2.303 0.0693 1 0.25 = 2.303 0.0693 0.602 = 20 min. (Also: t_ 75 % = 2 t_ 1/2 for first order.)
Second Order
For second-order reactions (Rate = k[A]²), integration gives 1/[A] = 1/[A₀] + kt. The reciprocal of concentration increases linearly with time. Plot of 1/[A] vs t gives a straight line with slope = k.
For a second-order reaction where Rate = k[A]^2 , integrating gives:
1 [A] = 1 [A]_0 + kt
k = 1 t ( 1 [A] - 1 [A]_0 ) Units: L mol^ -1 s^ -1
Plot of 1/[A] vs t → straight line, slope = +k , intercept = 1/[A]_0
Units of k : L mol ^ -1 s ^ -1 (or M ^ -1 s ^ -1 )
Second-order reactions slow down much faster than first-order as [A] drops
Two main experimental methods to find order: (1) Initial rate method — vary one reactant's concentration, measure initial rate, find the power relationship. (2) Graphical method — plot concentration data and see which integrated form gives a straight line.
There are two standard methods to determine the order of a reaction experimentally:
Initial Rate Method: Run multiple experiments, changing only ONE reactant's concentration each time. If doubling [A] doubles the rate → first order in A . If doubling [A] quadruples rate → second order in A
Graphical (Integrated Rate) Method: Take concentration vs time data from one experiment. Try plotting [A] vs t , [A] vs t , and 1/[A] vs t . Whichever gives a straight line tells you the order
Half-life Method: Measure t_ 1/2 at different initial concentrations. If t_ 1/2 is constant → first order. If t_ 1/2 [A]_0 → zero order. If t_ 1/2 1/[A]_0 → second order
From Exp 1→2: [A] doubles ( 0.1 0.2 ), rate quadruples ( 0.02 0.08 ). So 2^n = 4 n = 2 . Verify with Exp 1→3: [A] triples, rate increases 0.18/0.02 = 9 = 3^2 . Confirmed: second order.
Half-Life
Half-life (t₁/₂) is the time for reactant concentration to drop to half its initial value. Zero order: t₁/₂ = [A₀]/2k (depends on initial concentration). First order: t₁/₂ = 0.693/k (independent of initial concentration — the defining feature). Second order: t₁/₂ = 1/(k[A₀]). General relation: t₁/₂ ∝ [A₀]^(1−n) for nth order.
t_ 1/2 = 0.693 k (first order)
Zero-Order Half-Life
For zero-order reactions, t₁/₂ = [A₀]/2k. Half-life depends directly on the initial concentration — start with more, takes longer to reach half.
For a zero-order reaction: [A] = [A]_0 - kt . Setting [A] = [A]_0/2 :
t_ 1/2 = [A]_0 2k
t_ 1/2 [A]_0 — if you start with double the concentration, half-life doubles
Each successive half-life is shorter than the previous one (because the 'new [A]_0 ' after each half-life is smaller)
The reaction reaches completion in finite time: t_ complete = [A]_0/k = 2 t_ 1/2
This contrasts sharply with first-order, where the reaction theoretically never completes
First-Order Half-Life
For first-order reactions, t₁/₂ = 0.693/k = ln 2/k. It is independent of initial concentration. This is the most famous result in chemical kinetics — and identical to radioactive decay half-life.
For first order: [A]_0 [A] = kt . Setting [A] = [A]_0/2 : 2 = kt_ 1/2
t_ 1/2 = 2 k = 0.693 k
t_ 1/2 is independent of [A]_0 — the hallmark of first-order kinetics
After 1 half-life: 50% remains. After 2 half-lives: 25%. After 3: 12.5%. After n : (1/2)^n remains
[A] = [A]_0 ( 1 2 )^ t/t_ 1/2 — quick formula for any time
Time for 75% completion = 2t_ 1/2 . Time for 87.5% completion = 3t_ 1/2 . Time for 99.9% = 10t_ 1/2
If t_ 1/2 = 10 min, then k = 0.693/10 = 0.0693 min ^ -1
Quick recall
A first-order reaction is 87.5% complete. How many half-lives have passed?
For second-order reactions, t₁/₂ = 1/(k[A₀]). Half-life increases as concentration decreases — the reaction slows down dramatically as reactant is consumed.
For second order: 1 [A] = 1 [A]_0 + kt . Setting [A] = [A]_0/2 :
t_ 1/2 = 1 k[A]_0
t_ 1/2 1/[A]_0 — start with more concentrated solution → shorter half-life
Each successive half-life is longer than the previous (opposite of zero order!)
The reaction theoretically never completes (like first order), but slows down much faster
General Half-Life Relation
For an nth-order reaction: t₁/₂ ∝ [A₀]^(1−n). This single relation unifies all orders and is the key to determining order from half-life experiments.
t_ 1/2 [A]_0^ 1-n for nth order reaction
n = 0
t_ 1/2 [A]_0^ 1 → increases with [A]_0
n = 1
t_ 1/2 [A]_0^ 0 = constant → independent of [A]_0
n = 2
t_ 1/2 [A]_0^ -1 → decreases with [A]_0
n = 3
t_ 1/2 [A]_0^ -2 → decreases sharply with [A]_0
Quick recall
When [A]_0 = 0.1 M, t_ 1/2 = 100 s. When [A]_0 = 0.2 M, t_ 1/2 = 50 s. Find order.
The Arrhenius equation k = Ae^(−Eₐ/RT) shows how the rate constant increases exponentially with temperature. A is the pre-exponential (frequency) factor, Eₐ is the activation energy. The log form: log k = log A − Eₐ/(2.303RT). Plotting log k vs 1/T gives a straight line with slope = −Eₐ/(2.303R). For two temperatures: log(k₂/k₁) = (Eₐ/2.303R)(1/T₁ − 1/T₂).
k = A e^ -E_a/RT
Arrhenius Equation
The Arrhenius equation quantifies how the rate constant k changes with temperature T. The exponential dependence explains why even small temperature increases can dramatically speed up reactions.
Svante Arrhenius (1889) proposed that the rate constant depends on temperature as:
k = A e^ -E_a/RT
k = rate constant (what we want to find at any temperature)
A = pre-exponential factor (or frequency factor) — represents the frequency of collisions with correct orientation. Units same as k
E_a = activation energy (J mol ^ -1 ) — minimum energy required for reaction to occur
R = gas constant = 8.314 J mol ^ -1 K ^ -1
T = absolute temperature in kelvin (NEVER Celsius in this equation!)
e^ -E_a/RT = Boltzmann factor — fraction of molecules with energy E_a
Logarithmic Form
Taking the natural log of the Arrhenius equation gives ln k = ln A − Eₐ/RT. Plotting ln k (or log k) vs 1/T gives a straight line whose slope directly gives the activation energy.
Taking _ 10 of both sides of k = Ae^ -E_a/RT :
k = A - E_a 2.303 R T
This is in the form y = c + mx where y = k , x = 1/T , slope m = -E_a/(2.303R) , and intercept c = A .
Plot of k vs 1/T → straight line with slope = -E_a/(2.303R)
From slope: E_a = -slope 2.303 R = -slope 2.303 8.314 J mol ^ -1
Intercept on k axis = A (gives the frequency factor)
Steeper slope → higher E_a → reaction is more sensitive to temperature changes
This plot is called an Arrhenius plot — a standard tool in kinetics research
Two-Temperature Form
When k is known at two different temperatures T₁ and T₂, the activation energy can be found directly without graphing: log(k₂/k₁) = (Eₐ/2.303R)(1/T₁ − 1/T₂). This is the most tested Arrhenius formula in NEET.
Writing the Arrhenius log form at two temperatures and subtracting:
k_2 k_1 = E_a 2.303 R ( 1 T_1 - 1 T_2 )
k_1 at temperature T_1 (lower), k_2 at temperature T_2 (higher)
If T_2 > T_1 , then k_2 > k_1 , and (1/T_1 - 1/T_2) > 0 — both sides positive ✓
This eliminates the need for A — you don't need to know the frequency factor
Can also be used to find k at a new temperature if E_a and one k value are known
Quick recall
A reaction has k = 2 10^ -2 s ^ -1 at 300 K and k = 8 10^ -2 s ^ -1 at 320 K. Find E_a .
Activation energy (Eₐ) is the minimum extra energy reactant molecules must possess over their average energy to convert into products. Higher Eₐ → slower reaction. Eₐ can be lowered by a catalyst but not by changing temperature.
For a reaction to occur, colliding molecules must have energy ≥ a threshold energy ( E_ threshold ). The activation energy is the difference between this threshold and the average energy of the reactants:
E_a = E_ threshold - E_ reactants (average)
E_a is always positive (even for exothermic reactions — they still need an initial energy push)
Higher E_a → fewer molecules have enough energy → slower reaction
E_a for forward reaction ≠ E_a for backward reaction (unless H = 0 )
Relationship: E_a(forward) - E_a(backward) = H (enthalpy of reaction)
A catalyst lowers E_a by providing an alternative pathway — it does NOT change H
E_a
Decreases / Unchanged
H
Unchanged / Unchanged
k (rate constant)
Increases / Increases
Equilibrium constant K
Unchanged / Changes
Fraction above E_a
More molecules qualify / More molecules qualify
Collision Theory
Collision theory explains rates at the molecular level. For a reaction to occur, molecules must (1) collide, (2) with sufficient energy (≥ Eₐ), and (3) with proper orientation. Rate = Z_AB × e^(−Eₐ/RT) × p, where Z_AB is collision frequency and p is the steric (orientation) factor. The theory explains concentration and temperature effects but has limitations for complex reactions.
Rate = Z_ AB e^ -E_a/RT p
Three Conditions for Reaction
Not every collision leads to a reaction. Molecules must satisfy three simultaneous conditions: they must actually collide (frequency), the collision must have enough energy (threshold energy), and the molecules must be oriented correctly (steric factor).
According to collision theory, a chemical reaction occurs when reactant molecules collide. But not every collision is productive. Three conditions must be met simultaneously:
Condition 1 — Collision must occur: Molecules must physically meet. Collision frequency Z_ AB depends on concentration, molecular size, and temperature
Condition 2 — Sufficient energy: The kinetic energy along the line of centers must equal or exceed the threshold energy ( E_ threshold ). Only fraction e^ -E_a/RT of collisions satisfy this
Condition 3 — Proper orientation: Molecules must approach each other in a specific geometric arrangement. The steric factor p (always < 1) accounts for this
Collisions that satisfy ALL three conditions are called effective collisions (or fruitful collisions). The rate of reaction equals the rate of effective collisions.
Rate = Z_ AB e^ -E_a/RT p
Z_ AB = collision frequency (collisions per unit volume per unit time) — depends on concentration and temperature
e^ -E_a/RT = energy factor — fraction of collisions with energy ≥ E_a
p = steric factor (probability factor or orientation factor) — typically 10^ -1 to 10^ -9
Effective collisions = Z_ AB e^ -E_a/RT p — this is usually a tiny fraction of total collisions
Threshold Energy & Activation Energy
Threshold energy is the minimum total energy molecules must have for a collision to be effective. Activation energy = threshold energy − average kinetic energy of reactants. At higher temperatures, more molecules cross this threshold.
The threshold energy ( E_ threshold ) is the minimum energy a colliding pair must possess for bonds to break and new bonds to form. Not all molecules have this energy — only those in the high-energy tail of the Maxwell-Boltzmann distribution.
E_a = E_ threshold - E _ reactants
At temperature T , the fraction of molecules with energy E_a is given by the Boltzmann factor: f = e^ -E_a/RT .
At 300 K with E_a = 75 kJ/mol: f = e^ -75000/(8.314 300) e^ -30 10^ -13 — only 1 in 10^ 13 molecules has enough energy!
At 310 K: f 2 10^ -13 — doubles! This is why a 10°C rise roughly doubles the rate
Increasing temperature shifts the Maxwell-Boltzmann distribution → more molecules in the high-energy tail
A catalyst does NOT shift the distribution — it lowers the wall ( E_a ) so more existing molecules qualify
Steric Factor
The steric factor p accounts for the requirement that molecules must collide with the correct geometric orientation. For simple atoms p ≈ 1; for complex molecules p can be as small as 10⁻⁹. It explains why collision theory overestimates rates for complex molecules.
Even if two molecules collide with sufficient energy, the reaction won't happen if they approach from the wrong angle. The steric factor p (also called probability factor) corrects for this:
k = p Z_ AB e^ -E_a/RT
0 < p 1 — always a fraction (never greater than 1)
For reactions between atoms or simple molecules: p 1
For complex organic molecules: p can be as low as 10^ -9
NCERT example: For the reaction NO_2 + NO_2 arrow N_2O_4 , the N atoms must face each other — other orientations don't lead to product
The steric factor is the main reason collision theory overestimates rates for complex reactions
Limitations of Collision Theory
Collision theory works well for simple gas-phase bimolecular reactions but fails for: reactions in solution, unimolecular reactions, complex multi-step reactions, and reactions involving large molecules where p is hard to predict.
While collision theory provides an intuitive molecular-level picture of reaction rates, it has significant limitations:
Treats molecules as hard spheres — ignores actual molecular structure and electron clouds
Cannot predict the steric factor p — it must be determined experimentally for each reaction, making the theory semi-empirical
Fails for unimolecular reactions — how does a single molecule 'collide' with itself? (Lindemann mechanism explains this, but goes beyond simple collision theory)
Does not account for reactions in solution — solvent cage effects, diffusion, and solvation are ignored
Overestimates rates significantly for complex molecules because it cannot predict the correct orientation requirement
Temperature & Catalysis Effects
Temperature coefficient is the ratio of rate constants at temperatures 10°C apart: typically 2–3 for most reactions. A catalyst provides an alternative reaction pathway with lower activation energy, increasing rate without being consumed. Catalysts are classified as homogeneous (same phase) or heterogeneous (different phase). Enzyme catalysis follows Michaelis-Menten kinetics with substrate specificity.
Temperature coefficient = k_ T+10 k_T 2 to 3
Temperature Coefficient
The temperature coefficient is the factor by which the rate constant increases when temperature is raised by 10°C (or 10 K). For most reactions, this factor is between 2 and 3, meaning rate roughly doubles or triples with every 10°C rise.
The temperature coefficient ( ) of a reaction is defined as:
= k_ T+10 k_T 2 to 3
For most reactions, 2 – 3
If = 2 and temperature rises by 30°C: rate increases by 2^3 = 8 times
If = 3 and temperature rises by 20°C: rate increases by 3^2 = 9 times
General formula: rate increases by ^ T/10 for a temperature rise of T °C
This rule is approximate — the Arrhenius equation is the exact relationship
Quick recall
A reaction has temperature coefficient = 2. If the reaction takes 40 min at 300 K, how long at 320 K?
Temperature rise = 20 K = 2 steps of 10 K. Rate increases by 2^2 = 4 times. Time decreases by factor 4: t = 40/4 = 10 min.
Catalysis
A catalyst increases reaction rate by providing an alternative reaction pathway with lower activation energy. It participates in the reaction mechanism but is regenerated — not consumed. It does not change ΔH, ΔG, or the equilibrium constant.
A catalyst is a substance that increases the rate of a reaction without being consumed in the overall process. It works by providing an alternative mechanism with a lower activation energy ( E_a'< E_a ).
Catalyst lowers E_a → more molecules have sufficient energy → rate increases
Catalyst is not consumed — it is regenerated at the end of the reaction cycle
Catalyst does NOT change: H , G , S , or the equilibrium constant K
Catalyst speeds up BOTH forward and reverse reactions equally → equilibrium is reached faster but the position of equilibrium is unchanged
Catalyst does NOT appear in the overall balanced equation (though it appears in the mechanism steps)
A small amount of catalyst can catalyze a large amount of reaction (turnover)
Homogeneous vs Heterogeneous Catalysis
In homogeneous catalysis, the catalyst is in the same phase as reactants (e.g., acid catalysis in solution). In heterogeneous catalysis, the catalyst is in a different phase (usually solid catalyst with gaseous/liquid reactants). Heterogeneous catalysis involves adsorption on the catalyst surface.
Catalyst in SAME phase as reactants
Catalyst in DIFFERENT phase from reactants
Example: H⁺ catalyzing ester hydrolysis (all in solution)
Example: Fe in Haber process (solid Fe, gaseous N₂/H₂)
Example: NO catalyzing O₃ decomposition (all gases)
Example: Pt in contact process (solid Pt, gaseous SO₂/O₂)
Catalyst is uniformly mixed
Reaction occurs on catalyst SURFACE
Harder to separate and recover
Easy to separate (different phase)
Every catalyst molecule is accessible
Only surface atoms are active
Homogeneous examples for NEET: H ^+ in ester hydrolysis, NO in lead chamber process ( 2SO_2 + O_2 NO 2SO_3 )
Heterogeneous examples for NEET: Fe in Haber process ( N_2 + 3H_2 Fe 2NH_3 ), V _2 O _5 in contact process, Pt/Pd in catalytic converters
In heterogeneous catalysis, the steps are: (1) Diffusion to surface, (2) Adsorption on surface, (3) Reaction on surface, (4) Desorption of products, (5) Diffusion away
Adsorption is the key step — reactants are held on the catalyst surface, bonds are weakened, orientation is favorable
Enzyme Catalysis
Enzymes are biological catalysts (proteins) with extraordinary specificity and efficiency. They follow the lock-and-key model and Michaelis-Menten kinetics. At low substrate concentration, rate is first order; at high concentration, rate becomes zero order (enzyme saturation).
Enzymes are biological catalysts — proteins that catalyze biochemical reactions with remarkable speed and specificity. A single enzyme molecule can catalyze 10^3 to 10^7 reactions per second.
Lock and key model: Each enzyme has an active site with a specific shape. Only the correct substrate (reactant) fits — like a specific key in a specific lock
Specificity: Urease catalyzes ONLY the hydrolysis of urea — nothing else. This extreme selectivity is unique to enzymes
Optimal conditions: Each enzyme works best at a specific pH and temperature. Human enzymes typically work best at pH 7.4 and 37°C
Denaturation: At high temperatures or extreme pH, the enzyme's protein structure unfolds → it loses its shape and catalytic activity
Enzymes reduce E_a dramatically — sometimes by 50–75% compared to uncatalyzed reactions
Michaelis-Menten kinetics: At low [S] (substrate concentration), rate [S] (first order — plenty of free enzyme). At high [S] , all enzyme active sites are occupied → rate reaches a maximum ( V_ ) and becomes independent of [S] (zero order — enzyme saturation).