The Whole Chapter in One Glance

Chemical kinetics is about how fast reactions go. We measure rate, write the experimental rate law Rate=k[A]x[B]y\text{Rate} = k[A]^x[B]^y, classify reactions by order and molecularity, integrate the rate law for zero and first order, compute half-lives, and explain the steep temperature dependence with the Arrhenius equation, collision theory and catalysis.

Chemical kinetics chapter overview mind map

Slot every concept and numerical into one of these six branches.

Master Formula Sheet

Rate: Rate=−1ad[A]dt=+1cd[C]dt\text{Rate} = -\dfrac{1}{a}\dfrac{d[A]}{dt} = +\dfrac{1}{c}\dfrac{d[C]}{dt} (units mol L−1^{-1} s−1^{-1}).

Rate law: Rate=k[A]x[B]y\text{Rate} = k[A]^x[B]^y; order =x+y= x + y (experimental).

Zero order: [R]=[R]0−kt[R] = [R]_0 - kt; t1/2=[R]02kt_{1/2} = \dfrac{[R]_0}{2k}; kk in mol L−1^{-1} s−1^{-1}.

First order: k=2.303tlog⁡[R]0[R]k = \dfrac{2.303}{t}\log\dfrac{[R]_0}{[R]}; [R]=[R]0e−kt[R] = [R]_0 e^{-kt}; t1/2=0.693kt_{1/2} = \dfrac{0.693}{k}; kk in s−1^{-1}.

Arrhenius: k=Ae−Ea/RTk = A e^{-E_a/RT}; ln⁡k=ln⁡A−EaRT\ln k = \ln A - \dfrac{E_a}{RT}; log⁡k2k1=Ea2.303R(T2−T1T1T2)\log\dfrac{k_2}{k_1} = \dfrac{E_a}{2.303R}\left(\dfrac{T_2-T_1}{T_1 T_2}\right).

Units of kk: order 0 -> mol L−1^{-1} s−1^{-1}; order 1 -> s−1^{-1}; order 2 -> mol−1^{-1} L s−1^{-1}; order 3 -> mol−2^{-2} L2^2 s−1^{-1}.

Constants: R=8.314R = 8.314 J K−1^{-1} mol−1^{-1}; 0.693=ln⁡20.693 = \ln 2; 2.303=ln⁡102.303 = \ln 10.

Quick Comparison Tables

Order vs Molecularity

Feature Order Molecularity
Nature Experimental Theoretical
Values 0, fractional, negative allowed Positive integer (1-3)
Applies to Overall & elementary Elementary steps only

Zero vs First Order

Feature Zero order First order
Rate law Rate =k= k Rate =k[R]= k[R]
Integrated [R]=[R]0−kt[R] = [R]_0 - kt k=2.303tlog⁡[R]0[R]k = \frac{2.303}{t}\log\frac{[R]_0}{[R]}
Units of kk mol L−1^{-1} s−1^{-1} s−1^{-1}
t1/2t_{1/2} [R]0/2k[R]_0/2k (depends on [R]0[R]_0) 0.693/k0.693/k (independent of [R]0[R]_0)
Straight-line plot [R][R] vs tt log⁡[R]\log[R] vs tt

Last-Minute Memory Hooks

  • Order = experimental (read off the rate law, found by experiment); molecularity = mechanistic (integer, elementary steps only). For an elementary step, they are usually the same.
  • Units of kk tell the order: mol L−1^{-1} s−1^{-1} -> zero; s−1^{-1} -> first; L mol−1^{-1} s−1^{-1} -> second.
  • First-order t1/2=0.693/kt_{1/2} = 0.693/k is independent of [R]0[R]_0; zero-order t1/2=[R]0/2kt_{1/2} = [R]_0/2k depends on [R]0[R]_0.
  • After nn half-lives a first-order reactant is down to (1/2)n(1/2)^n.
  • Pseudo-first-order: look for one reactant in large excess, so its concentration remains effectively constant.
  • Arrhenius: plot ln⁡k\ln k vs 1/T1/T -> slope −Ea/R-E_a/R. Rate roughly doubles per 10 K.
  • Catalyst lowers EaE_a, speeds forward and reverse reactions equally, and does NOT change ΔH\Delta H or KK.
  • Effective collision needs energy ≥Ea\ge E_a AND proper orientation.

One-line revision flow: Rate -> rate law -> order/molecularity -> zero/first order integrated -> half-life -> Arrhenius/temperature -> collision theory & catalyst.