Chemical Kinetics mind map

Chemical Kinetics: what the chapter covers

The main idea of each topic with its first key points. Open a branch above for everything, or read the full notes.

1. 1Chapter Overview

   Chemical Kinetics studies how fast chemical reactions occur and why their speeds change under different conditions. In NCERT and NEET, the chapter revolves around rate of reaction, rate law, order, molecularity, integrated rate equations, half-life, Arrhenius equation, activation energy and collision theory. Thermodynamics tells whether a reaction is possible, but kinetics tells how quickly it reaches completion. A reaction may be thermodynamically feasible but kinetically slow, like conversion of diamond to graphite. This chapter is highly numerical and graph-based. NEET commonly asks units of rate constant, graphical identification of order, half-life dependence, pseudo first order reactions and activation energy from Arrhenius plots.

   • The chapter has two major numerical areas: integrated rate laws and Arrhenius equation.
   • Graph slopes are very important: zero order [R] vs t has slope -k, first order log[R] vs t has slope -k/2.303, and Arrhenius plot log k vs 1/T has slope -Ea/2.303R.
   • Order may be zero, fractional, integral or negative; molecularity is always a positive integer.
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2. 2Rate of Reaction

   Rate of reaction tells how quickly reactants are consumed or products are formed. For a reaction, it is expressed as change in concentration per unit time and usually has unit mol L⁻¹ s⁻¹. Average rate is calculated over a finite time interval, while instantaneous rate is the slope of tangent to concentration-time curve at a particular moment. Rate expression must consider stoichiometric coefficients so that the same reaction rate is obtained from any reactant or product. Rate law expresses rate in terms of molar concentrations raised to experimentally determined powers. NEET often tests signs, units, stoichiometric division, rate law interpretation and numerical substitution.

   • For aA + bB → cC + dD, rate = -1/a d[A]/dt = -1/b d[B]/dt = 1/c d[C]/dt = 1/d d[D]/dt.
   • Instantaneous rate is more accurate for changing rates and is obtained graphically from tangent slope.
   • Rate law is not the same as balanced chemical equation for complex reactions.
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3. 3Factors Affecting Rate

   Reaction rate changes because the number and effectiveness of molecular collisions change. Ionic reactions in aqueous solution are often very fast, while covalent bond-breaking reactions may be slower because old bonds must break and new bonds must form. Increasing concentration or pressure usually increases collision frequency. Raising temperature increases kinetic energy and greatly increases the fraction of molecules crossing activation energy. A catalyst provides an alternate pathway with lower activation energy and speeds both forward and reverse reactions. Greater surface area allows more particles to be exposed for reaction. NEET frequently asks qualitative effects, catalyst misconceptions, temperature coefficient and graphical observations.

   • Rate increases when effective collision frequency increases.
   • Temperature affects both kinetic energy and Arrhenius factor e^(-Ea/RT).
   • A catalyst does not change ΔG, ΔH or equilibrium constant; it only helps reach equilibrium faster.
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4. 4Order & Molecularity

   Order and molecularity describe reaction dependence and mechanism, but they are not the same. Order is the sum of powers of concentration terms in the experimentally determined rate law. It may be zero, fractional, integral or even negative. Molecularity is the number of reacting species participating in one elementary step and is always a positive integer. Zero order reactions have rate independent of reactant concentration, first order reactions have rate proportional to one concentration term, and second order reactions have total concentration power two. Pseudo first order reactions are actually higher order but behave as first order because one reactant is in large excess, such as hydrolysis of ester in water.

   • Overall order = sum of powers in rate law.
   • Molecularity indicates how many species collide in an elementary event.
   • Complex reactions occur through multiple elementary steps, and the slowest step often controls rate.
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5. 5Integrated Rate Equations

   Integrated rate equations connect reactant concentration with time and allow calculation of rate constant, remaining concentration and half-life. They are essential because experimental concentration-time data is easier to measure than instantaneous rate. For zero order reactions, concentration falls linearly with time and half-life is directly proportional to initial concentration. For first order reactions, logarithmic concentration changes linearly with time and half-life is constant. For second order reactions, reciprocal concentration varies linearly with time and half-life is inversely proportional to initial concentration. NEET repeatedly asks identification of order from straight-line graph, slope formulas, half-life patterns and first order numerical calculations using log values.

   • Integrated rate laws are used when concentration changes continuously with time.
   • For first order reactions, equal time intervals reduce concentration by the same fraction.
   • In first order decay, after n half-lives, remaining amount = initial amount/2^n.
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6. 6Arrhenius Equation & Collision Theory

   Arrhenius equation explains the strong dependence of rate constant on temperature and activation energy. Molecules must acquire minimum energy to form an activated complex or transition state before converting into products. This minimum extra energy is activation energy. A small rise in temperature can cause a large increase in rate because many more molecules cross the activation energy barrier. Collision theory adds that not every collision gives product; collisions must have sufficient energy and proper orientation. The Arrhenius plot of log k versus 1/T is a straight line whose slope gives activation energy. NEET commonly asks activation energy calculation, catalyst effect and effective collision concepts.

   • Threshold energy is the minimum total energy needed for reaction.
   • Activation energy = threshold energy - average energy of reactants.
   • Only a fraction of collisions are effective.
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