Chemical Kinetics studies how fast reactions go and what controls their speed, covering rate of reaction, rate law, order and molecularity, and the integrated rate equations for zero and first order reactions. It also explains half-life, the Arrhenius equation linking rate to temperature and activation energy, and how collision theory and catalysts work. It is a high-yield, formula and numerical heavy chapter that NEET tests almost every year.
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Rate of Reaction
The rate of a reaction is how fast a reactant disappears or a product forms per unit time.
Units: mol L⁻1 s⁻1. Reactants get a minus sign, products a plus sign.
- Average rate = Δ[X]/Δt; instantaneous rate = slope of conc-time curve
- Divide each species by its coefficient to get one common rate
- For N2 + 3H2 → 2NH3, H2 vanishes 3× and NH3 forms 2× as fast as N2
Rate Law
The rate law links rate to reactant concentrations and is found only by experiment.
k = rate constant (depends on temperature, not on concentration).
- Exponents x, y are experimental, NOT the balanced-equation coefficients
- For multistep reactions the slowest (rate-determining) step fixes the rate law
- k changes with temperature; doubling [A] that doubles rate ⇒ order 1 in A
Order vs Molecularity
Order is the experimental concentration dependence; molecularity counts colliding species in an elementary step.
Order can be 0, fractional or negative; molecularity is always a positive integer.
- Order comes from experiment; molecularity comes from the mechanism
- Molecularity is never zero or fractional
- Higher order means more sensitive to concentration, not necessarily faster
Units of the Rate Constant
The units of k depend on the overall order, so units alone reveal the order.
Because rate is always mol L⁻1 s⁻1, the units of k must balance it.
- Zero order: k in mol L⁻1 s⁻1
- First order: k in s⁻1
- Second order: k in L mol⁻1 s⁻1
Zero-Order Integrated Rate
A zero-order reaction proceeds at a fixed rate independent of concentration.
Plot of [A] vs t is a straight line with slope −k.
- Rate is constant no matter how much reactant remains
- Examples: decomposition of NH3 on hot Pt, many photochemical reactions
- Slope of [A]-vs-t line gives −k directly
First-Order Integrated Rate
For first order the rate is proportional to the remaining reactant, so it slows as it proceeds.
Plot of log[A] vs t is linear with slope −k/2.303. The 2.303 converts ln to log10.
- All radioactive decay is first order
- k is independent of initial concentration and of concentration units
- Use fraction REMAINING (not % reacted) in [A]0/[A]
Half-Life (t½) and Order
Half-life is the time for the reactant concentration to fall to half its value.
0.693 = ln 2. General rule: t½ ∝ [A]01⁻ⁿ.
- First-order t½ is independent of initial concentration (its hallmark)
- Zero-order t½ is directly proportional to [A]0
- After n half-lives, fraction remaining = (1/2)ⁿ
Arrhenius Equation
Rate rises steeply with temperature because more molecules cross the activation barrier.
T in Kelvin; rule of thumb: rate roughly doubles per 10°C rise.
- Plot of log k vs 1/T is a straight line of slope −Eₐ/(2.303 R)
- Two-temperature form: log(k2/k1) = (Eₐ/2.303R)·[(T2−T1)/(T1T2)]
- Lower Eₐ ⇒ faster reaction; heating shifts Maxwell-Boltzmann to higher energy
Collision Theory
A collision leads to reaction only if it has enough energy and the correct orientation.
Z_AB = collision frequency; P = steric (orientation) factor, usually < 1.
- Two conditions: energy ≥ Eₐ AND proper molecular orientation
- e^(−Eₐ/RT) is the fraction of energetic (effective) collisions
- P corrects collision theory, which otherwise overpredicts the rate
Catalysis
A catalyst speeds a reaction by offering an alternative path of lower activation energy.
It is not consumed; it is regenerated at the end.
- Does not change ΔH or the position of equilibrium
- Lowers Eₐ for both directions, so equilibrium is reached faster
- A small drop in Eₐ gives a huge rise in rate (Eₐ sits in an exponent)
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Frequently Asked Questions
The rate of a reaction is the change in concentration of a reactant or product per unit time, with units of mol per litre per second. For a reaction aA giving cC it is written as minus one over a times d[A]/dt, equal to plus one over c times d[C]/dt, so a single common rate value is obtained.
Order is the sum of the powers of concentration in the experimentally found rate law and can be zero, fractional, or a whole number. Molecularity is the number of species colliding in a single elementary step and is always a positive integer, never zero or fractional.
For a first order reaction the half-life is t-half equal to 0.693 divided by k, where 0.693 is the natural log of 2. A key feature is that this half-life is independent of the initial concentration, which is a hallmark of first order reactions, and all radioactive decay is first order.
Yes, Chemical Kinetics is part of the NEET Class 12 Chemistry syllabus and is a reliable scoring chapter that usually contributes one to two questions every year. The questions are mostly numerical, based on rate law, order, integrated rate equations, half-life, and the Arrhenius equation.
Raising the temperature increases the rate sharply, and as a rule of thumb the rate roughly doubles for every 10 degree Celsius rise. This is captured by the Arrhenius equation, k equals A times e to the power minus Ea over RT, where A is the frequency factor, Ea is the activation energy, R is the gas constant, and T is the temperature in Kelvin.