Unit 7: Reaction Kinetics — Long Questions

Ans. Reaction Kinetics

Definition Reaction kinetics is the study of the rates of chemical reactions and the factors that affect the rates of chemical reactions.

The studies of reaction kinetics involves:
• A variety of experimental methods, measuring reaction rates, determining reaction orders, understanding reaction mechanisms.
• It is a common observation that the rates of different chemical reactions vary greatly.

Types of reaction rates

Fast reaction The reaction of NaCl with AgNO₃ is very fast.

Moderate reaction The hydrolysis of an ester occurs at a moderate rate.

Slow reaction The rusting of iron is a slow process.

Importance of reaction rates in industry

The rate of a chemical reaction and its control are often crucial in industrial processes. These factors can:
• Determine the economic feasibility of using a particular reaction on a commercial scale.
• Affect production efficiency, safety, and cost-effectiveness.

COLLISION THEORY

Collision theory

Collision theory explains how and why chemical reactions occur at the molecular level.

Basic requirements for a reaction

According to the theory • Reactant particles (atoms, ions, or molecules) must form a homogeneous mixture.
• They must collide with each other for a reaction to take place.

Collision depends upon energy of colliding molecule. It can be:
• Effective – leading to product formation
• Ineffective – particles bounce back unchanged

Conditions for effective collisions

For a collision to be effective, two main conditions must be met:

1. Sufficient Energy: The colliding particles must possess at least a minimum amount of energy known as the activation energy (Eₐ).

2. Proper Orientation: The particles must approach each other in a specific orientation that allows bond formation or breaking.

Activation Energy The minimum energy required for an effective collision between the reacting specie is called activation energy.

Most reactions are slow

• Not all collisions are effective, many particles may lack the required energy or proper orientation.
• This explains why most chemical reactions are slow and require specific conditions to proceed efficiently.

Ans. Rate of reaction

Definition The rate of a chemical reaction is defined as "The change in concentration of a reactant or a product divide by the time taken for the change."

This can be mathematically expressed as:
Rate =Δ[Concentration]/ ΔTime or
Rate of reaction = Δx / Δt

Δx is the change in concentration of a reactant or product.
Δt is the time interval over which the change occurs

Change in concentration during a reaction

During a chemical reaction, the reactants are converted into products. As the reaction proceeds, the concentration of reactants decreases because they are being consumed.
The concentration of products increases as they are being formed.

Graphical explanation

Consider the irreversible reaction for a reactant A which is changing irreversibly to B.
A → B
• The concentration of reactant A decreases over time.
• The concentration of product B increases over time.
• The slope is steepest at the beginning, indicating a high reaction rate.
• Rapid decrease in concentration of reactants and rapid increase in concentration of products.
• As time passes, the slope gradually flattens, showing the rate is decreasing.
• Eventually, the curves become almost horizontal parallel, meaning:
• No further change in concentrations
• The reaction has completed

The rate of reaction is not constant — it changes at every moment during the reaction.

Units of reaction rate

The rate of a chemical reaction has the units of concentration by time. Concentration is expressed in moldm⁻³ and the time is expressed in seconds so, the units for the rate of reaction are moldm⁻³ s⁻¹.

Rate of reaction = moles dm⁻³ / seconds = mol.dm⁻³ s⁻¹

However, for slower reactions, the time unit may be in minutes or hours, giving units like mol dm⁻³ min⁻¹ or mol dm⁻³ h⁻¹.

In gas-phase reactions, pressure is often used instead of concentration, with rate units like atm/s or kPa/min.

Reaction rate

For a general reaction A→B
The rate of reaction can be written in terms of either:
• The rate of disappearance of reactant A
• The rate of appearance of product B

Mathematically
Rate of reaction = -Δ[A]/Δt = +Δ[B]/Δt

The negative sign indicates that the concentration of A is decreasing over time.
The positive sign for B reflects that its concentration is increasing.

Average rate and instantaneous rate of reaction

Average rate of reaction

The rate of reaction between two specific time intervals or the rate over time period.
It gives a general idea of how fast a reaction occurs between two points in time.

Instantaneous Rate of Reaction

The rate at any one instant during the intervals is called instantaneous rate.
Instantaneous Rate= d [Reactant or Product]/dt

Relationship between the average and instantaneous rate: The average and instantaneous rates are equal only at one specific point during the interval.
At the beginning of the reaction, the instantaneous rate is usually greater than the average rate because concentrations are higher, which typically speeds up the reaction.
At the end of the interval, the instantaneous rate tends to be lower than the average rate as reactant concentrations drop.

Ans. To measure the rate of a chemical reaction, scientists monitor how the concentration of reactants or products change with respect to time at regular intervals as reaction proceeds.

Decomposition of HI

1. Taking Samples at Regular Time Intervals
Small samples are withdrawn from the reaction mixture at set time intervals.

2. Determining Concentration
The concentration of a reactant (which decreases) or a product (which increases) is measured for each sample.

3. Plotting Data
A graph is plotted between time on x-axis and concentration on y-axis.
Average rate (over a time interval), and
Instantaneous rate (at a specific time point) can be calculated.

Let us consider the decomposition of hydrogen iodide (HI) into hydrogen (H₂) and iodine (I₂) at 508°C.
experimental data for decomposition of HI is given in table 7.1.

2HI₍g₎ → H₂₍g₎ + I₂₍g₎

The following data shows how the concentration of HI decreases over time

A graph is plotted by taking
X-axis: Time (in seconds)
Y-axis: Concentration of HI (in mol dm⁻³)

HI is a reactant, its concentration decreases over time, resulting in a falling curve.
• The steepness (slope) of the curve indicates the progress of the reaction.
• A steeper slope near the beginning implies a faster rate.
The slope becomes less steep as the reaction slows down.

Determining Instantaneous Rate To find the reaction rate (e.g., 100 seconds):
(i) Draw a tangent to the curve at 100 seconds.
(ii) Form a right-angled triangle with the tangent line. The slope of the tangent is the rate of reaction after 100 sec. at that point.
• A right angled triangle ABC is completed with the tangent as hypotenuse.
• Suppose the change in concentration (ΔC) is 0.04 mol dm⁻³
• The time interval (Δt) is 100 seconds
(iii) Use the formula
Rate of reaction=ΔC/Δt = 0.04 mol dm⁻³/100 s = 4×10⁻⁴ mol dm⁻³ s⁻¹
This means the concentration of HI is decreasing at a rate of 4 × 10⁻⁴ mol dm⁻³ s⁻¹ at 100 seconds.

Product Concentration Graph If we instead plot the concentration of products (H₂ or I₂) versus time:
• The curve would rise, since product concentrations increase.
• The slope of the tangent at 100 seconds would also give the same numerical value for the rate:
4×10⁻⁴ mol dm⁻³ s⁻¹

Ans. The change in concentrations of reactants or products can be determined by both physical and chemical methods depending upon the type of reactants or products involved.

(a) Chemical Method

Hydrolysis of ester

This method is particularly suitable for reactions in solution, where it is possible to chemically analyze a reactant or product at different times. A typical example is the acid-catalyzed hydrolysis of an ester, ethyl acetate. In the presence of small amount of an acid.

Reaction
CH₃COOC₂H₅₍aq₎ + H₂O₍l₎ → CH₃COOH₍aq₎ + C₂H₅OH₍aq₎

Here, ethyl acetate reacts with water in the presence of an acid catalyst to produce acetic acid and ethanol.

Procedure

(i) Start the Reaction: Mix ethyl acetate, water, and a small amount of acid as a catalyst.
(ii) Sampling: At fixed time intervals, withdraw a small sample using a pipette.
(iii) Quenching: Immediately transfer the sample into about four times its volume of ice-cold water.
• This dilutes and chills the mixture, effectively stopping the reaction at that moment.
(iv) Titration: Titrate the quenched sample against a standard NaOH solution using phenolphthalein as an indicator
• This neutralizes the acetic acid formed, allowing its concentration to be determined.

Purpose By performing this titration at various time intervals, the concentration of acetic acid formed can be measured.
This provides data about the concentration of acetic acid formed during the reaction at different time intervals.

(b) Physical Methods

Some of the methods used for the measurement of concentration are as follows:

(i) Spectrophotometry or colorimetry

This method is applicable when a reactant or product absorbs light in the ultraviolet (UV), visible, or infrared (IR) radiations.
The rate of reaction can be measured by measuring the amount of radiation absorbed. For the reaction shown in figure, the concentration can be measured using the calorimetry.

(ii) Electrical conductivity method

Conductometric method for measuring reaction rate

This method is particularly useful for studying reactions that involve ionic species, either as reactants or products.
The electrical conductivity of an aqueous solution depends on the concentration and mobility of ions formed during the reaction.
As the reaction proceeds, the number and type of ions change, leading to a measurable change in conductivity.
The rate of change in conductivity reflects the rate of change in ion concentration, which in turn relates to the rate of the reaction.

(iii) Volume change method

Gas Volume Measurement Method for Reaction Rates

This method is useful for reactions where there is a change in the volume of gas — either due to the production or consumption of a gaseous substance.
During a reaction, if gas is produced or consumed, the total gas volume in the system changes.
This change in volume is directly proportional to the extent of reaction and the change in concentration of gaseous reactants or products.
By measuring gas volume at regular time intervals, the rate of reaction can be determined.

Ans. Factors affecting the rate of chemical reaction

The rates at which reactants are consumed and products are formed during chemical reactions vary greatly. Even a chemical reaction involving the same reactants may have different rates under different conditions.

The factors affecting the rates of reactions are

(i) Concentrations of the reactants
(ii) Temperature of the system
(iii) Surface area
(iv) Catalyst

(i) Concentration

According to the law of mass action
"The rate of a chemical reaction is directly proportional to the product of the concentrations of the reactants, each raised to a power equal to its coefficient in the balanced chemical equation."

• Higher concentration means more particles per unit volume.
• This leads to increased collision frequency between reactant molecules.
• More collisions → greater chance of effective collisions → faster reaction.

(ii) Temperature

Increase in temperature increases, the reaction rate.

(Maxwell-Boltzmann distribution curve)

Effect of Temperature on Reaction Rate

Molecules collide more frequently and more energetically
A larger fraction of molecules has energy equal to or greater than activation energy (Eₐ).
This results in a higher rate of effective collisions, hence a faster reaction rate.
The rate roughly doubles or triples with every 10 °C rise in temperature.
Higher temperature = more molecules with sufficient energy (Eₐ) → increased reaction rate.

Boltzmann distribution curve

The curve shows the distribution of kinetic energies among particles at a given temperature.
Most particles have moderate energy, some have low, and some have high energy.
Eₐ (activation energy) is a threshold: only particles with energy equal to Eₐ can react upon collision.
Shaded area under the curve beyond Eₐ represents particles that can successfully react.
As temperature increases, the curve flattens and shifts right, increasing the shaded area.

Raising the temperature increases the average kinetic energy of particles.
Particles move faster, leading to:
• More frequent collisions.
• Higher-energy collisions.

A greater proportion of particles now have energy equal to activation energy (Eₐ).
On the Boltzmann distribution curve:
• The peak flattens and shifts to the right.
• The shaded area beyond Eₐ (successful collisions) approximately doubles for each 10 °C rise in temperature.

Conclusion Increasing the temperature increases the rate of reaction.

(iii) Catalyst

Definition A catalyst is a substance which alters the rate of a chemical reaction but remains chemically unchanged at the end of the reaction.
Catalysts are usually present in small amounts compared to reactants.

Examples

(i) Platinum (Pt) speeds up the reaction between H₂ and O₂ to form water.
(ii) Manganese dioxide (MnO₂) speeds up the decomposition of KClO₃.
(iii) Copper(II) chloride (CuCl₂) catalyzes the oxidation of HCl to Cl₂.

Energy profile with catalyst

• The presence of a catalyst lowers the energy barrier, as shown in the energy profile diagram.
• The overall energy change (ΔH) of the reaction remains the same, but the activation energy (Eₐ) is lowered.
• This results in faster reactions because more molecules can overcome the new, lower activation energy.

Ans. Catalysis

Definition The process which take place in the presence of a catalyst is called catalysis.

Types of Catalysis

(i) Homogeneous catalysis

In this process, the catalyst and the reactants are in the same phase and the reacting system is homogeneous throughout. The catalyst is distributed uniformly throughout the system.

Example

The formation of SO₃₍g₎ from SO₂₍g₎ and O₂₍g₎ in the contact process for the manufacture of sulphuric acid, needs NO₍g₎ as a catalyst. Both the reactants and the catalyst are gases.

2SO₂₍g₎ + O₂₍g₎ → NO₍g₎ → 2SO₃₍g₎

Esters are hydrolyzed in the presence of H₂SO₄. Both the reactants and the catalyst are in the solution state.

CH₃COOCH₃₍ₗ₎ + H₂O₍ₗ₎ → H₂SO₄ → CH₃COOH₍aq₎ + CH₃CH₂OH₍aq₎

(ii) Heterogeneous Catalysis

Definition The process in which the catalyst and the reactants are in different phases is called heterogeneous catalysis mostly, the catalysts are in the solid phase.
The reactants are in the gaseous or liquid phase.

Example Oxidation of ammonia to NO in the presence of platinum gauze to produce HNO₃ industrially through the heterogeneous catalysis.

4NH₃₍g₎ + 5O₂₍g₎ → Pt,800°C → 4NO₍g₎ + 6H₂O₍g₎

Hydrogenation of unsaturated organic compounds are catalysed by finely divided Ni, Pd or Pt. Ni at 150°C.

C₂H₄₍g₎ + H₂₍g₎ → Ni,150°C → C₂H₆₍g₎

Ans. Rate Law and Rate Constant
The rate of a chemical reaction at a given temperature may depend on the concentration of one or more reactants and products.
Definition of rate law: The representation of rate of a reaction in terms of concentration of the reactants is known as rate law.
A rate law is an equation that relates the rate of a reaction to the concentrations of reactants raised to various powers according to the experimental data.
General reaction: The reaction between A and B where 'a' moles of A and 'b' moles of B react to form 'c' moles of C and 'd' moles of D.

aA + bB → cC + dD

We can write the rate equation as

Rate = k [A]ˣ [B]ʸ

Where x and y are the experimentally determined values that may or may not be equal to the coefficient of reactants in the balanced chemical equation, as 'a' and 'b' in the above equation. This expression is called rate equation. The brackets [ ] represent the molar concentrations and the proportionality constant k is called rate constant for the reaction.
If [A] = 1 mol dm⁻³ and [B] = 1 mol dm⁻³
Rate of reaction = k × 1ˣ × 1ʸ = k (rate constant)
Definition of rate constant: The rate constant can be defined as "The specific rate constant of a chemical reaction is the rate of reaction when the concentration of the reactants are unity". Under the given conditions, k remains constant, but it changes with temperature.

Ans. Reaction Order
Definition: "The order of a reaction with respect to a specific reactant is the exponent applied to that reactant's concentration within the rate equation".
Consider a general reaction

aA +bB → cC + dD
Rate = k [A]ˣ[B]ʸ

• The exponents 'x' and 'y' in the above equation give the order of reaction with respect to the individual reactants.
• Thus, the reaction is of order 'x' with respect to A and of order 'y' with respect to B. The overall order of reaction is (x+y).
• The order of a reaction defines how the reactant concentration influences its rate. For a single-reactant, the order is simply the concentration's power in the rate equation.
• The chemical reactions are classified as zero, first, second and third order reactions. The order of reaction provides valuable information about the mechanism of a reaction.
• It is crucial to differentiate between the order concerning a single reactant and the overall reaction order.

Measurement of reaction order
Take equation for the reaction of nitrogen (II) oxide (NO) with H₂:

2H₂(g) + 2NO(g) → N₂(g) + 2H₂O(g)

The experimentally determined rate expression for this is,

Rate = k [H₂][NO]²

The expression shows that this reaction is:
(i) First-order with respect to H₂
(ii) Second-order with respect to NO
(iii) Third-order overall (1+2=3)

Types of reaction order
(i) Zero order reaction
Definition: The rate of a zero order reaction is independent of the concentration of the reactants.
Rate law

Rate=k[A]⁰=k

Since anything raised to the power of 0 is 1, the rate = constant. (k)
Chang in the concentration of reactants does not change the reaction rate.
Examples of zero order reactions
(i) H₂(g) + Cl₂(g) → 2HCl(g) (in presence of light):

• A photochemical reaction, where the rate is controlled by light intensity, not concentration.

(ii) 2NH₃(g) → N₂(g) + 3H₂(g) (on hot platinum surface):

• Surface-catalyzed reaction; rate depends on the availability of surface sites, not [NH₃].
• Photochemical reactions usually zero order.
• Often seen in photochemical and enzyme-catalyzed reactions at saturation.

(ii) First Order Reaction
Definition: A reaction for which some of the exponents of the concentrations in the rate equation is 1.
Rate law

Rate=k[A]¹=k[A]

This means doubling [A] doubles the rate, tripling [A] triples the rate, and so on.
In these reactions, there may be multiple reactants present but concentration of only one reactant affects the rate, even if others are present.
The reaction follows exponential decay of the reactant over time.
Common in decomposition and radioactive decay processes.
Example
Decomposition of nitrogen pentoxide is an example of first order as it rate of reaction depends upon one molecules.

2N₂O₅(g) → 2N₂O₄(g) + O₂(g)

Rate law

Rate=k[N₂O₅]

This equation suggests that the reaction is first order.
Overall order = 1
All radio-active disintegration reactions are first order reactions.

(iii) Second order reactions
Definition
A reaction for which sum of the exponents of the concentrations in the rate equation is 2. A second order is that reaction whose rate depends upon the concentration of one reactant raised to the second power or on the concentrations of the two different reactants, each raised to the first power. Hypothetically, it can be expressed as:
Rate laws
There are two common forms:
(a) Single reactant

Rate=k[A]²

(b) Two different reactants

Rate=k[A][B]

• Doubling [A] in Rate=k[A]² increases rate by a factor of 4.
• In Rate=k[A][B], doubling both A and B increases rate by a factor of 4.

Units of k

mol⁻¹·dm³·s⁻¹

Oxidation of nitric oxide Oxidization of nitric acid with ozone is first order with respect to NO and first order with respect to O₃. The sum of individual orders gives the overall order of reaction.

NO(g) + O₃(g) → NO₂(g) + O₂(g)

Rate law

Rate=k[NO][O₃]

First order with respect to NO, first order with respect to O₃, overall order = 2.

(iv) Third-Order Reactions
Definition: A third order reaction is the reaction for which sum of the exponents of the concentration in the rate equation is 3.
Forms of Rate Law
Single reactant

Rate=k[A]³

Two reactants

Rate=k[A]²[B]

Three reactants

Rate=k[A][B][C]

Examples The reaction involves eight reactant molecules but experimentally it has been found to be a third order reaction.

2FeCl₃(aq) + 6KI(aq) → 2FeI₂(aq) + 6KCl(aq) + I₂

Experimental rate law

Rate=k[FeCl₃][KI]²

Order with respect to FeCl₃ = 1, order with respect to KI = 2, overall order = 3.

2NO(g)+O₂(g)→2NO₂(g)

Rate law

Rate=k[NO]²[O₂]

Overall order = 3

(v) Fractional Order Reactions
Definition: A reaction in which sum of the exponents of rate equation is in fraction, is called the fractional order reaction.
Example: Consider the formation of carbon tetrachloride.

CHCl₃(g) + Cl₂(g) → CCl₄(g) + HCl(g)

Rate law

Rate=k[CHCl₃]¹[Cl₂]^(1/2)

Order with respect to CHCl₃ = 1
Order with respect to Cl₂ = ½
Overall order = 1 + ½ = 1.5
Reactions involving free radicals frequently exhibit fractional orders.
Fractional orders often indicate that the reaction mechanism is complex, typically involving free radicals.
In this reaction Cl₂ dissociates into Cl• radicals, which then initiate the reaction with CHCl₃. Because the formation of Cl• is in dynamic equilibrium, the concentration of Cl₂ affects the rate non-linearly, leading to the ½ order.

Ans. Units of rate constant
The rate constant is specific for a particular reaction at a certain temperature.
Reason: Concentrations are in mol dm⁻³ and the reaction rate is in units of The units for k depend on the order of the reaction and the units of time. when mole mol dm⁻³ s⁻¹ the order of reaction, the number of molecules at which rate actually depends will charge so, on it of rate constant will also change.
General equation:

Rate = k[Reactants]ⁿ where n = order of reaction

k = Rate/(Reactants)ⁿ = (moldm⁻³)s⁻¹/(moldm⁻³)ⁿ

kₙ = (mol. dm⁻³)¹⁻ⁿ s⁻¹
kₙ = (concentration)¹⁻ⁿs⁻¹

This equation can be used to determine units of any order of reaction.

(i) Units of k of zero order
For a zero order reaction (n = 0),

k₀ = (mol. dm⁻³)¹⁻⁰s⁻¹
k₀ = mol. dm⁻³s⁻¹

(ii) Units of k for first order
For a first order reaction (n = 1), the rate is directly proportional to the concentration of one reactant.

k₁ = (mol. dm⁻³)¹⁻¹ s⁻¹
k₁ = (mol. dm⁻³)⁰s⁻¹
k₁ = s⁻¹

Therefore, the units of k for a first order rate constant are s⁻¹.

(iii) Units of k for second order
For a second order reaction (n = 2),

k₂ = (mol. dm⁻³)¹⁻² s⁻¹
k₂ = (mol⁻¹. dm³)⁻¹ s⁻¹
k₂ = dm³ mol⁻¹s⁻¹

The units of k for a second order rate constant are dm³ mol⁻¹ s⁻¹.

(iv) Units of k for third order
For a third order reaction (n = 3),

k₃ = (mol. dm⁻³)¹⁻³ s⁻¹
k₃ = mol⁻² dm⁶ s⁻¹
k₃ = dm⁶ mol⁻² s⁻¹

Therefore, the units of k for a third order rate constant are dm⁶ mol⁻² s⁻¹.

Ans. Determination of rate constant
The rate constant (k) of a reaction can be calculated using the following two methods:
(i) Initial Concentration Method
Reaction of H₂O₂ and I⁻ ion
In the presence of hydrogen ions, hydrogen peroxide H₂O₂, reacts with iodide ions to form water and iodine.

H₂O₂(aq) + 2I⁻(aq) + 2H⁺(aq) → 2H₂O(l) + I₂(aq)

The rate equation for this reaction is

rate of reaction = k[H₂O₂] [I⁻]

The progress of the reaction can be followed by measuring the initial rate of formation of iodine. Table 7.2 shows the rates of reaction obtained using various initial concentrations of each reactant.
Calculation of rate constant
The procedure for calculating k is shown below, using the data for experiment 1.
Step 1: Write out the rate equation

rate of reaction = k[H₂O₂] [I⁻]

Step 2 Rearranged the equation in terms of k

k = rate/[H₂O₂][I⁻]

Step 3 Substitute the values

k = 3.50×10⁻⁶/(0.0200)×(0.0100)
k = 1.75 × 10⁻² dm³ mol⁻¹ s⁻¹

Table 7.2 Effect of change in concentrations of reactants on the rate of reaction.
Experiment: [H₂O₂] moldm⁻³: [I⁻] moldm⁻³: [H⁺] moldm⁻³: Initial rate of reaction moldm⁻³ s⁻¹
1: 0.0200: 0.0100: 0.0100: 3.5 × 10⁻⁶
2: 0.0300: 0.0100: 0.0100: 5.3 × 10⁻⁶
3: 0.0050: 0.0200: 0.0200: 1.75 × 10⁻⁶

The concentration of hydrogen ions is ignored because [H⁺] does not appear in the rate equation. The reaction is zero order with respect to [H⁺].

(ii) Half-life method
Definition: Half-Life (t₁/₂) is "the time taken for the concentration of a reactant to fall to half of its original value"
Measuring reaction order

• Rate constant (k) is calculated by measuring half life time it takes for the concentration of reactant to decrease by half.
• For first-order reactions, the half-life is independent of concentration and follows a fixed relationship.
• For a first-order reaction, the rate constant and half-life are related by:

k = 0.693 ÷ t₁/₂

Decomposition of hydrogen peroxide
Given

• The half-life of hydrogen peroxide is 2 hours
• Reaction is first-order
• To calculate k, we first need to convert the half-life, which is 2 hours, into seconds

2 × 60 × 60 = 7200 s

Put this value into the equation

k = 0.693/7200 = 9.6×10⁻³s⁻¹

Ans. Reaction mechanism
A reaction mechanism is a detailed, step-by-step description of how a chemical reaction occurs at the molecular level to yield the products.
Role of reaction mechanism

• The overall balanced equation shows only: Reactants and products.
• The reaction mechanism shows: Actual steps (called elementary steps) through which reactants turn into products
• Each elementary step represents a single molecular event, like: Bond breaking, bond forming and most reactions proceed through multiple steps, not just one. Each step is triggered by the collision of: Atoms, ions and molecules. Molecularity: The number of reactant molecules involved in an elementary step. Types of Molecularity (i) Unimolecular Reaction: A unimolecular elementary reaction involves only a single reactant molecule. Example: Decomposition of N₂O₅.

N₂O₅(g) → NO₂(g) + NO₃(g)

(ii) Bimolecular Reaction: An elementary reaction involves two atoms, ions or molecules and is called bimolecular
Examples

CO(g) + NO₃(g) → NO₂(g) + CO₂(g)
NO(g) + O₃(g) → NO₂(g) + O₂(g)

(iii) Termolecular Reaction: A termolecular reaction involves the simultaneous reaction of three molecules.

• Involves three reacting molecules simultaneously
• Rare due to the low probability of three particles colliding at once Example: formation of ozone in the atmosphere by reaction of oxygen molecule and atomic oxygen in smog formation.

2O₂(g) + O(g) → O₃(g) + O₂(g)

Intermediates
Definition: They are short lived species that are produced mechanism and consumed in a subsequent step. They are not stable so they do not mansion in balanced chemical equation
Example: carbocation (CH₃)₃ C⁺ etc.

Rate-determining step
Definition: In many reactions, one step is significantly slower than all of the others. This step is called rate-determining step.

• Controls the overall rate of the reaction. Balanced equation is net result of sum of all individual steps.
• Any step after the Rate determining steps does not affect the rate.
• Only reactants involved in the Rate determining steps appear in the rate law.
• All the reactants appear in rate determining step will also appear in the rate equation.
• Determines the order of the reaction.
• Even if the overall reaction involves several steps, the rate-determining step acts like a bottleneck.
• The rate equation is written based only on the rate-determining step.
• This helps chemists predict or control the speed of complex reactions. Example: The overall reaction between NO and H₂ is given.

2NO(g) + 2H₂(g) → 2H₂O(g) + N₂(g)

Then the rate law would be:

Rate=k[NO]²[H₂]¹

Even though 2NO and H₂ are used in the overall reaction, only one of each appears in the rate law because they are involved in the rate determining step.

Table 7.3 Effect of change in concentrations of reactants on the rate of reaction
[NO] in (mol dm⁻³): [H₂] in (mol dm⁻³): Initial rate (atm min⁻¹)
0.006: 0.001: 0.025
0.006: 0.002: 0.050
0.006: 0.003: 0.0075
0.001: 0.009: 0.0063
0.002: 0.009: 0.025
0.003: 0.009: 0.056

Experimental setup
Six experiments were conducted to determine how the rate of reaction depends on the concentration of H₂ and NO.
The general form of the reaction: Rate∝[H₂]ᵐ[NO]ʸ

Effect of H₂ Concentration
In these experiments, [NO] is kept constant, and only [H₂] is varied.
Doubling [H₂] then Rate doubles
Tripling [H₂] the Rate triples
Rate∝[H₂]¹
The reaction is first order with respect to H₂.

Effect of NO Concentration
Here, [H₂] is constant, and [NO] is varied.
Doubling [NO] results Rate increases 4 times
Tripling [NO] results Rate increases 9 times
Conclusion

Rate∝ [NO]²

The reaction is second order with respect to NO
The overall rate equation of reaction is,

Rate ∝ [H₂][NO]²
or Rate = k[H₂]¹[NO]²

The reaction is a third order one.
This final equation is the rate law for this reaction.
It should be kept in mind that rate law cannot be predicted from the balanced chemical equation.

The possible mechanism consisting of two steps for the reaction is as follows.
Step 1: 2NO(g) + H₂(g) → N₂(g) + H₂O₂(g) (rate determining step)
Step 2: H₂O₂(g) + H₂(g) → 2H₂O(g) (Fast)

Reaction of F₂ and NO₂
The step 1 is slow and rate determining.
The reaction between nitrogen dioxide and fluorine gas:

2NO₂(g) + F₂(g) → 2NO₂F(g)

This reaction is first order in NO₂, first order in F₂ and second order overall. The experimental rate law is first order in NO₂ and in F₂:

Rate = k[NO₂][F₂] (Observed)

The accepted mechanism for the reaction is:
NO₂(g) + F₂(g) → NO₂F(g) + F(g) Rate = k₁[NO₂][F₂]
NO₂(g) + F(g) → NO₂F(g) Rate = k₂[NO₂][ F⁻ ]

The first step is slow and determines the rate, in agreement with the observed rate expression. The second and fast step does not affect the reaction rate because fluorine atoms are produced and react with NO₂ as soon as they are produced.