Order vs
Molecularity

Comparison of order and molecularity

The correct statement

Order is experimental and unrestricted; molecularity counts colliding particles and must be a small whole number.

Order = sum of exponents ; Molecularity = number of colliding species

1. Order comes from experiment and can take any value the data supports — zero, fractional or negative.
2. Molecularity counts actual particles taking part in a collision, and you cannot have zero or half a particle.
3. So order can be zero or fractional while molecularity cannot.
4. NCERT states this as conclusion (i) after the H₂O₂ mechanism.

A ten-particle equation with experimental order 2

The valid conclusion

A large mismatch between apparent and actual order proves a multi-step mechanism.

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1. Counting reactant particles gives 1 + 6 + 3 = 10, suggesting tenth order if it were elementary.
2. The measured order is 2, so the reaction cannot be elementary.
3. Ten particles colliding simultaneously is effectively impossible, confirming a multi-step path.
4. Since the reaction is complex, molecularity has no meaning overall, so (c) is wrong.

Step 1: H₂O₂ + I⁻ → H₂O + IO⁻ (slow); Step 2: H₂O₂ + IO⁻ → H₂O + I⁻ + O₂ (fast)

The intermediate

An intermediate is produced in one step and consumed in a later one, and never appears in the overall equation.

Overall: 2H₂O₂ → 2H₂O + O₂

1. IO⁻ is produced in step 1 and consumed in step 2.
2. It does not appear in the overall balanced equation.
3. That is precisely the definition of an intermediate.
4. I⁻ is consumed first and regenerated later, making it the catalyst rather than the intermediate.

I⁻ is consumed in step 1 and regenerated in step 2

Role of I⁻

A catalyst is consumed early and returned later, emerging unchanged overall.

—

1. I⁻ is used up in the slow first step.
2. It reappears in the fast second step.
3. Overall it is neither created nor destroyed, and it speeds the reaction up.
4. That makes it a catalyst, not an intermediate.

Four reactions

The unimolecular one

Unimolecular means a single species reacting on its own.

—

1. NH₄NO₂ decomposes by itself with no collision partner needed — one reacting species.
2. 2HI involves two HI molecules colliding, so bimolecular.
3. 2NO + O₂ involves three particles, so termolecular.
4. H₂ + I₂ involves two, so bimolecular.

Termolecular reactions

Reason for their rarity

Simultaneous three-body encounters are statistically improbable.

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1. A collision requires particles to arrive at the same place at the same instant with suitable orientation.
2. Arranging that for two particles is common; for three it is far less likely.
3. NCERT states directly that the probability of more than three molecules colliding and reacting simultaneously is very small.
4. So termolecular reactions are rare and slow, and molecularity above three is effectively impossible.

A complex reaction with a rate-determining step

Relation between slow-step molecularity and overall order

The slow step controls the rate, so its molecularity sets the observed order.

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1. The overall rate is governed entirely by the slowest step.
2. The rate law therefore reflects the species involved in that step.
3. So the number of species colliding in the slow step equals the overall order.
4. This is NCERT's conclusion (iii) stated verbatim.

Applicability of molecularity

Where it is meaningless

Molecularity describes a single elementary step, so it cannot describe a multi-step process.

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1. Molecularity counts species colliding in one elementary reaction.
2. A complex reaction consists of several elementary steps, each with its own molecularity.
3. There is no single number that describes the whole sequence.
4. So for a complex reaction, overall molecularity has no meaning — NCERT conclusion (ii).

An elementary reaction

Relation between order and molecularity

In a single-step reaction, the molecules that must collide are exactly those appearing in the rate law.

For an elementary step, exponents = coefficients of that step

1. An elementary reaction occurs in one step, so the equation IS the mechanism.
2. The molecules that must collide are precisely those whose concentrations control the rate.
3. So the exponents in the rate law equal the coefficients, making order equal molecularity.
4. NCERT's Summary states this directly, and it is only valid for elementary reactions.

Order = 1.5

What this implies

No single collision can produce a fractional exponent, so a fraction proves a mechanism.

Molecularity ∈ {1, 2, 3} only

1. Molecularity counts particles and must be a whole number, so 1.5 is impossible as a molecularity.
2. If the reaction were elementary, order would equal molecularity and would therefore be a whole number.
3. Since the measured order is fractional, the reaction cannot be elementary.
4. It must proceed through several steps, typically involving a dissociation pre-equilibrium.

Four quantities

Which cannot be zero

Molecularity counts particles, and a reaction with zero particles is meaningless.

—

1. Order can be zero, as in ammonia decomposition on platinum.
2. The rate constant is never exactly zero in practice, but activation energy can be zero for a barrierless reaction.
3. Molecularity counts the species colliding in an elementary step, and that count must be at least 1.
4. So molecularity can never be zero.

Rate = k[NO]²[O₂], assumed elementary

Molecularity

For an elementary reaction, molecularity equals the total number of colliding species.

molecularity = sum of coefficients in the elementary step

1. Two NO molecules and one O₂ molecule must collide simultaneously.
2. Total colliding species = 2 + 1 = 3.
3. So the reaction is termolecular.
4. Consistently, the order is 2 + 1 = 3, matching the molecularity as expected for an elementary reaction.

A multi-step mechanism

Definition of the rate-determining step

A sequence can proceed no faster than its slowest link.

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1. In a sequence of steps, the overall pace is set by the slowest one.
2. NCERT compares this to a relay team, whose chances depend on the slowest runner.
3. The slow step has the highest activation energy, not the lowest, so (d) is wrong.
4. The rate law contains the species of the slow step.

Comparison of intermediate and catalyst

The distinguishing feature

Both are absent from the overall equation; only the sequence of events differs.

—

1. An intermediate does not exist at the start; it appears during the reaction and is then used up.
2. A catalyst is present at the start, is consumed, and reappears by the end.
3. Neither appears in the overall balanced equation, which is why they are confused.
4. So the distinguishing feature is purely the order of formation and consumption.

Types of reaction listed by NCERT

What they exemplify

Any reaction not completed in a single step is complex.

—

1. NCERT lists consecutive reactions (such as oxidation of ethane through alcohol, aldehyde and acid), reverse reactions and side reactions.
2. All of these involve more than one elementary step.
3. Reactions completed in one step are elementary; sequences of such steps are complex.
4. So all three named types are complex reactions.

Definition question

Meaning of bimolecular

The prefix counts the colliding species.

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1. Bi- means two.
2. So a bimolecular elementary reaction requires two species to collide simultaneously.
3. NCERT's example is the dissociation of hydrogen iodide, 2HI → H₂ + I₂.

A complex reaction

The true statement

The slowest step is the bottleneck and therefore fixes the observed order.

—

1. A complex reaction proceeds through a sequence of elementary steps, so (d) is wrong.
2. Molecularity has no overall meaning for a complex reaction, so (c) is wrong.
3. Order cannot be predicted from stoichiometry, as the KClO₃ example proves, so (b) is wrong.
4. The overall rate is limited by the slowest step, which therefore determines the order.

Two-step mechanism with the first step slow

Rate law

The rate law contains the species of the rate-determining step.

Rate law reflects the slow step

1. The first step is slow and therefore rate-determining.
2. It involves one A and one B colliding.
3. So Rate = k[A][B], first order in each and second order overall.
4. The second B, consumed in the fast step, does not enter the rate law — this is exactly why order differs from overall stoichiometry.

A + 2B → C with Rate = k[A]

Interpretation

Species absent from the rate law take part only after the slow step.

Rate law reflects the slow step only

1. B appears in the balanced equation, so it certainly reacts — option (b) is wrong.
2. Its absence from the rate law means changing [B] does not affect the rate.
3. That happens when B enters only in a fast step, after the bottleneck has been passed.
4. So B is not involved in the rate-determining step.

Four quantities

The theoretical one

Molecularity is deduced from a proposed mechanism, not measured.

—

1. Order, rate constant and half-life are all obtained from experimental data.
2. Molecularity is read off a proposed elementary step and cannot be measured directly.
3. It is a theoretical concept describing what is believed to happen at the molecular level.

Ethane oxidation via several intermediates

Type of reaction

A sequence of steps producing successive intermediates is a consecutive complex reaction.

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1. The reaction passes through a series of distinguishable intermediate species.
2. Each conversion is a separate elementary step.
3. A sequence of elementary steps constitutes a complex reaction.
4. NCERT cites this exact example when introducing consecutive reactions.

Nitration of phenol giving two products

What it illustrates

Competing pathways from the same reactant are side reactions.

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1. Two different products form from the same starting material by competing routes.
2. Such competing pathways are called side reactions.
3. NCERT lists side reactions among the types of complex reaction, alongside consecutive and reverse reactions.

Elementary reaction with order 3

Molecularity

For elementary reactions the two are equal.

order = molecularity for an elementary step

1. The reaction is stated to be elementary.
2. For an elementary reaction, order equals molecularity.
3. So the molecularity is 3, making it termolecular.
4. Consistency check: 3 is a permitted molecularity, though such reactions are rare.

Candidate values

Impossible molecularities

Molecularity is restricted to 1, 2 and 3.

molecularity ∈ {1, 2, 3}

1. Zero is impossible because at least one species must react.
2. One half is impossible because particles cannot be fractional.
3. Four is effectively impossible because simultaneous four-body collisions do not occur.
4. Only 1, 2 and 3 are permitted, so option (a) correctly identifies all three impossibilities.

Step 1: H₂O₂ + I⁻ → H₂O + IO⁻ ; Step 2: H₂O₂ + IO⁻ → H₂O + I⁻ + O₂

Nature of the two steps

Count the colliding species in each step.

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1. Step 1 involves H₂O₂ colliding with I⁻ — two species.
2. Step 2 involves H₂O₂ colliding with IO⁻ — again two species.
3. Both are single steps involving two colliding species, hence bimolecular elementary reactions.
4. NCERT states this explicitly.

Definition of order with respect to a reactant

Where the power is read from

Order is defined by the rate law, and only the rate law.

Rate = k[A]ˣ[B]ʸ

1. NCERT defines order with respect to a reactant as the power of its concentration appearing in the rate law equation.
2. The balanced equation gives stoichiometry, not order.
3. The equilibrium constant expression describes the position of equilibrium, not the rate.
4. The Arrhenius equation concerns temperature dependence and contains no concentration at all.

A single bimolecular elementary step

Overall order

For an elementary step, order equals molecularity.

order = molecularity = 2

1. Bimolecular means two species collide, so the molecularity is 2.
2. For an elementary reaction the order equals the molecularity.
3. So the reaction is second order overall.
4. It could be first order in each of two reactants, or second order in a single one.

The two-step mechanism and Rate = k[H₂O₂][I⁻]

The invalid conclusion

The overall equation contains no I⁻ at all, so it could not have predicted the rate law.

Overall: 2H₂O₂ → 2H₂O + O₂ ; Rate = k[H₂O₂][I⁻]

1. The overall equation is 2H₂O₂ → 2H₂O + O₂, which contains no iodide.
2. Yet the rate law depends on [I⁻], so the equation could not have predicted it.
3. Statements (b), (c) and (d) are all correct features of the mechanism.
4. So (a) is the invalid conclusion — and it is precisely the point NCERT uses this example to make.

A complex reaction whose first step is slow

What the intermediate's rate of formation determines

Nothing downstream can proceed faster than the bottleneck supplies material.

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1. The slow first step produces the intermediate.
2. Subsequent fast steps can only consume the intermediate as fast as it is supplied.
3. So the rate of the whole sequence is limited by how fast the intermediate forms.
4. NCERT states that the rate of formation of the intermediate will determine the rate of this reaction.