Chemical Kinetics NCERT Solutions help students understand how fast or slow chemical reactions occur and what factors influence their rates. This chapter is vital for board exams and competitive exams like JEE and NEET because it connects theoretical chemistry with practical applications. It teaches you how to express rate laws mathematically, interpret reaction mechanisms, and use graphs for experimental data analysis.
Chemical kinetics doesn’t just deal with how fast reactions go—it explains why they proceed at those rates. Students explore the collision theory, activation energy, and order of reaction through detailed examples. The chapter uses logical reasoning and calculation-based questions that strengthen both analytical and mathematical skills in chemistry.
In this post, you’ll find a complete breakdown of the Class 12 Chemistry NCERT chapter “Chemical Kinetics” — including definitions, derivations, important formulas, and solved conceptual explanations. Each section below is formatted for easy reading with step-by-step tables, so you can build strong conceptual clarity.
Chemical Kinetics NCERT Solutions
- Rate of Reaction and Basic Definitions
- Order and Molecularity of Reaction
- Rate Law and Rate Constant
- Factors Affecting Reaction Rate
- FAQs
Rate of Reaction and Basic Definitions
Understanding Average and Instantaneous Rate
| Type of Rate | Definition | Formula | Example |
|---|---|---|---|
| Average Rate | Change in concentration over a given time interval | \(ext{Rate}_{avg} = -\frac{\Delta [A]}{Delta t} = frac{Delta [B]}{Delta t}\) | Decomposition of H2O2 |
| Instantaneous Rate | Rate at a particular moment of time | \(ext{Rate}_{inst} = lim_{\Delta t o 0} \frac{Delta [A]}{Delta t}\) | Used in graphical calculations |
The rate of a chemical reaction quantifies how quickly reactants are converted into products. The rate can be expressed in terms of change in concentration of either reactants or products per unit time. For example, for the reaction \(A ightarrow B\), the rate decreases as the reaction proceeds because the concentration of A decreases over time. Instantaneous rate is found by drawing a tangent to the concentration–time graph.
Units of rate depend on the reaction order. Typically, the rate has units of \(mol,L^{-1},s^{-1}\). Understanding the difference between average and instantaneous rates helps students interpret experimental data accurately and predict reaction progress graphs effectively.
Order and Molecularity of Reaction
Comparing Order and Molecularity
| Property | Order | Molecularity |
|---|---|---|
| Definition | Sum of powers of concentration in rate law | Number of reacting species colliding |
| Determination | Experimental | Theoretical (mechanistic) |
| Possible Values | Can be fractional or zero | Always a whole number |
| Example | \(ext{Rate} = k[A]^1[B]^2 \Rightarrow ext{Order} = 3\) | Bimolecular reaction like \(2NO + O_2 ightarrow 2NO_2\) |
Order and molecularity seem similar but are conceptually different. The order of a reaction is determined experimentally from the rate law, while molecularity refers to the number of molecules colliding in a single step of the mechanism. The order can be fractional (for complex reactions) or zero, whereas molecularity cannot.
Students should focus on understanding that order provides quantitative data about how the rate depends on concentration. For example, in a first-order reaction, if the concentration doubles, the rate doubles as well. These relationships are often tested in numerical questions, and recognizing the difference between these two terms is crucial for board and entrance exam success.
Rate Law and Rate Constant
Integrated Rate Laws for Zero, First, and Second Order Reactions
| Order | Rate Law Expression | Integrated Form | Graph |
|---|---|---|---|
| Zero Order | \(Rate = k\) | \([A] = [A]_0 – kt\) | Linear decrease in concentration |
| First Order | \(Rate = k[A]\) | \(\ln[A]_0/[A] = kt\) | Straight line when ln[A] vs t is plotted |
| Second Order | \(Rate = k[A]^2\) | \(\frac{1}{[A]} – \frac{1}{[A]_0} = kt\) | Inverse relationship |
The integrated rate law helps in calculating the concentration of reactants at any time or determining the order of reaction from experimental data. The rate constant \(k\) is independent of concentration but depends on temperature and catalyst presence. For first-order reactions, plotting \(\ln[A]\) versus time gives a straight line, which is a common method in practical experiments.
Understanding these mathematical expressions helps you predict half-life and the time required for partial completion of reactions. For instance, the half-life for first-order reactions is given by \(t_{1/2} = \frac{0.693}{k}\), which remains constant irrespective of initial concentration. This concept is highly scoring and must be practiced thoroughly with numerical examples.
Factors Affecting Reaction Rate
Physical and Chemical Influences on Rate
| Factor | Effect on Rate | Explanation |
|---|---|---|
| Concentration | Increases with higher reactant concentration | More collisions occur per unit time |
| Temperature | Rate doubles for every 10°C rise (approx.) | Increases kinetic energy and collision frequency |
| Catalyst | Speeds up reaction | Lowers activation energy |
| Surface Area | Increases for heterogeneous reactions | Provides more reaction sites |
| Pressure | Increases in gaseous reactions | Higher collision probability |
The rate of a chemical reaction depends on several factors such as concentration, temperature, catalysts, and surface area. Among these, temperature has a particularly strong effect described by the Arrhenius equation: \(k = A e^{-E_a/RT}\), where \(E_a\) is the activation energy. The exponential relationship explains how small changes in temperature can cause large changes in rate constants.
Catalysts, on the other hand, offer an alternate reaction pathway with lower activation energy, effectively increasing the number of successful collisions. Understanding these dependencies enables students to apply kinetics to real-life contexts such as enzyme-catalyzed biological reactions, industrial synthesis, and chemical equilibrium adjustments.