# Chemical kinetics **Chemical kinetics**, also called reaction kinetics, is the branch of physical [[Chemistry|chemistry]] that measures how fast a [[Chemical_reaction|chemical reaction]] runs and explains the speed in terms of concentrations, [[Temperature|temperature]], catalysts and the sequence of molecular steps behind the overall equation. Its central object is the [[Rate_equation|rate law]], `rate = k[A]^m[B]^n`, in which the orders m and n are found by experiment rather than read from the balanced equation and the rate constant k depends on temperature.[^af-ratelaw] Integrating the rate law gives concentration against time, and from that form comes the [[Half-life|half-life]]. In the microsim below the reader slides the starting concentration [A]₀ from 0.01 to 2 mol/L across three reactions tuned to share a ten-minute half-life at 0.200 mol/L, and the three formulas `t_half = [A]0/(2k)`, `0.693/k` and `1/(k[A]0)` answer why only the first-order half-life refuses to move.[^af-halflives] On the Chemistry flagship's spine this page is the main article for Part VI — Reaction, section *Reaction rate and rate laws*, and its half-life sim is the shared embed (C11) that the Materials science and Energy flagships reuse where they meet first-order decay. Sibling pages carry the story on: [[Arrhenius_equation|the Arrhenius equation]] for temperature, [[Catalysis|catalysis]] for lowered barriers, [[Steady_state_(chemistry)|the steady state]] for mechanisms, and [[Michaelis–Menten_kinetics|Michaelis–Menten kinetics]] for enzymes. Kinetics is the complement of [[Chemical_thermodynamics|chemical thermodynamics]]: thermodynamics says whether a reaction can go and how far, kinetics says how long it takes and by what path. The two meet at [[Chemical_equilibrium|equilibrium]], where the forward and reverse rates are equal and the [[Equilibrium_constant|equilibrium constant]] is the ratio of the two rate constants.[^af-keq] ## History The first quantitative rate study is usually credited to Ludwig Wilhelmy, who followed the inversion of cane sugar in acid in 1850, and the law of mass action that ties rate to concentration was stated by Cato Guldberg and Peter Waage in 1864.[citation needed] The subject became a discipline with Jacobus van 't Hoff's work on chemical dynamics; the first Nobel Prize in Chemistry, in 1901, was awarded to him "in recognition of the extraordinary services he has rendered by the discovery of the laws of chemical dynamics and osmotic pressure in solutions".[^nobel-1901] Svante Arrhenius, whose exponential temperature law bears his name, received the prize in 1903 for his theory of electrolytic dissociation,[^nobel-1903] and Wilhelm Ostwald received it in 1909 "for his work on catalysis and for his investigations into the fundamental principles governing chemical equilibria and rates of reaction".[^nobel-1909] The twentieth century added the machinery for mechanisms: [[Collision_theory|collision theory]], the [[Transition_state|transition state]] and the [[Activated_complex|activated complex]], the steady-state treatment of [[Reaction_intermediate|intermediates]], and the [[Enzyme_kinetics|enzyme kinetics]] of Michaelis and Menten. Reactions too fast for a burette came within reach through flash photolysis and relaxation methods, for which Manfred Eigen, Ronald Norrish and George Porter shared the 1967 Nobel Prize in Chemistry "for their studies of extremely fast chemical reactions, effected by disturbing the equilibrium by means of very short pulses of energy".[^nobel-1967] ## Factors affecting reaction rate The Portal Book's kinetics chapter groups the levers that change a rate into the nature of the reactants, their physical state and surface area, their concentrations, the temperature, and the presence of a catalyst; for gases, [[Pressure|pressure]] stands in for concentration, and for photochemical reactions the absorbed light is a further lever.[^af-factors] Each is treated below, and the concentration lever carries this article's microsim. ### Nature of the reactants Some reactions are fast because the bonds that must break are weak or none must break at all; ionic precipitations in water are effectively instantaneous, while a covalent rearrangement such as the decomposition of cyclobutane to ethylene proceeds with a first-order rate constant of 9.2 × 10⁻³ s⁻¹ in the book's example, so that four-fifths of it has decomposed after about three minutes.[^af-cyclobutane] The identity of the reactants also fixes the [[Activation_energy|activation energy]], the barrier that the temperature lever works against.[^af-ea] ### Physical state Reactants in one phase, whether a [[Solution_(chemistry)|solution]] or a gas mixture, meet by molecular collision throughout the volume. When a reactant is a solid or an immiscible liquid the reaction is confined to the interface, and stirring, grinding or dispersing the phase changes the rate without changing the chemistry; the kinetics chapter treats physical state as its own factor for this reason.[^af-factors] ### Surface area of solid state For a reaction on a solid, the available surface is the effective concentration of the solid. The book's zero-order example is the decomposition of [[Ammonia|ammonia]] on hot [[Tungsten|tungsten]]: once the metal surface is saturated, adding more ammonia cannot speed the reaction, and the concentration falls linearly, from 0.0028 mol/L to 0.0001 mol/L in 35 minutes, with a half-life of 18 minutes.[^af-zero] Zero order of this kind is conditional, or "pseudo", behaviour that holds only while the surface stays saturated.[^af-zero-cond] The same logic underlies [[Heterogeneous_catalysis|heterogeneous catalysis]] and the study of reactions in [[Surface_science|surface science]]. ### Concentration The rate law `rate = k[A]^m[B]^n` states how the rate depends on the [[Molar_concentration|molar concentrations]] of the reactants. The orders m and n must be measured, and the units of k follow from the overall order: mol L⁻¹ s⁻¹ for zero order, s⁻¹ for first order and L mol⁻¹ s⁻¹ for second order.[^af-ratelaw] Integrating the rate law for a single reactant gives three families of curves. Zero order: `[A] = [A]0 − k·t`, with `t_half = [A]0/(2k)`. First order: `ln[A] = ln[A]0 − k·t`, or `[A] = [A]0·exp(−k·t)`, with `t_half = 0.693/k`. Second order: `1/[A] = 1/[A]0 + k·t`, with `t_half = 1/(k·[A]0)`.[^af-halflives][^boyd-integrated] The microsim on this page is the half-life test. Three reactions are tuned so that all three have a half-life of exactly 10 minutes when [A]₀ = 0.200 mol/L: a zero-order reaction with k₀ = 0.010 mol L⁻¹ min⁻¹, a first-order reaction with k₁ = 0.0693 min⁻¹ and a second-order reaction with k₂ = 0.50 L mol⁻¹ min⁻¹. These constants are ILLUSTRATIVE, chosen for the display rather than measured. The reader's one control is [A]₀ on a logarithmic slider from 0.01 to 2 mol/L, and the three curves `[A]0 − k·t`, `[A]0·exp(−k·t)` and `1/(1/[A]0 + k·t)` are redrawn with half-life staircase markers. Pushed to 2 mol/L, the zero-order half-life grows to 100 minutes and the second-order half-life shrinks to 1.0 minute while the first-order half-life stays at 10 minutes; pulled down to 0.020 mol/L the pattern reverses, 1.0 minute for zero order and 100 minutes for second order. The zero-order line is clamped at [A] = 0 once t exceeds [A]₀/k, because a reactant cannot go negative. The book's own numbers behave the same way. The dimerization of butadiene is second order with k = 5.76 × 10⁻² L mol⁻¹ min⁻¹; starting at 0.200 mol/L, 10.0 minutes leaves 0.179 mol/L, and the half-life at that start is about 87 minutes, a value that doubles if the starting concentration is halved.[^af-butadiene] Iodine-131 decays with a first-order constant of 0.138 d⁻¹ and a half-life of 5.02 days however much is present, which is why a fixed half-life is the signature of [[Radioactive_decay|radioactive decay]] and of every other first-order process.[^af-i131] When one reactant is in large excess its concentration barely changes and the reaction shows a pseudo-order with `k_obs = k·[B]0^n`, the book's route to measuring one order at a time.[^boyd-pseudo] ### Temperature Raising the temperature increases the fraction of collisions that carry enough energy to cross the activation barrier, and the rate constant rises through the [[Arrhenius_equation|Arrhenius equation]], `k = A·exp(−Ea/(R·T))`, with R = 8.314 J mol⁻¹ K⁻¹.[^af-arrhenius] The classroom rule that rates roughly double for every 10 °C rise is a rule of thumb, not a law; it holds near room temperature for barriers of the order of 50 kJ/mol and fails elsewhere.[^af-rule] The Portal Book's hydrogen iodide data run from 3.52 × 10⁻⁷ L mol⁻¹ s⁻¹ at 555 K to 3.95 × 10⁻² L mol⁻¹ s⁻¹ at 781 K, a factor of about 10⁵ over 226 K, and give an activation energy of about 1.8 × 10⁵ J/mol; the sibling article walks that line point by point.[^af-hi] ### Catalysts A catalyst offers a different [[Reaction_mechanism|mechanism]] with a lower activation energy for the rate-determining step while leaving the energies of reactants and products where they were, so the rate rises but the [[Equilibrium_constant|equilibrium constant]] does not.[^af-catalyst] In the book's two-path example the first step's barrier drops from 80 kJ to 70 kJ on the catalysed path while the second step's 20 kJ barrier is unchanged.[^af-paths] Because `K = kf/kr`, the catalyst must speed the forward and reverse reactions by the same factor; the [[Catalysis|catalysis]] article and its sim make that point with a barrier the reader can lower. ### Pressure For reactions among gases, concentration and [[Partial_pressure|partial pressure]] are the same lever, since by the [[Ideal_gas_law|ideal gas law]] the concentration n/V of each gas equals its partial pressure divided by RT. Compressing a gas mixture therefore raises every reactant concentration at once and speeds a bimolecular gas reaction roughly as the square of the compression. Pressure is also the variable most often followed in gas-phase kinetics, because a change in the number of moles of gas shows up as a change in total pressure at fixed volume. ### Absorption of light A [[Photon|photon]] can supply the energy that thermal collisions cannot, exciting a molecule into a state from which it dissociates or reacts. The rate of a photochemical step is set by the rate at which photons are absorbed rather than by temperature, which is why the [[Ozone–oxygen_cycle|ozone–oxygen cycle]] of the upper atmosphere and the light reactions of [[Photosynthesis|photosynthesis]] are driven by sunlight rather than by heat. Flash photolysis, described below, turns the same effect into an instrument for starting a reaction with a pulse. ## Experimental methods A rate is measured by following a concentration in time. The Portal Book's first-order example is the decomposition of hydrogen peroxide, whose concentration falls from 1.000 to 0.500, 0.250, 0.125 and 0.0625 mol/L at 6-hour intervals; a plot of ln[H₂O₂] against time is a straight line, which is the test for first order, and the slope gives the rate constant.[^af-h2o2] Concentrations are followed by [[Spectrophotometry|spectrophotometry]] where a species absorbs light, by [[Titration|titration]] of samples quenched at known times, by pressure for gas reactions and by conductivity for reactions that make or consume ions. Two analyses turn data into a rate law. The method of initial rates compares the starting rates of runs with different starting concentrations and reads each order from the ratio; for nitric oxide reacting with ozone, five trials give first order in each reactant.[^af-initial] The graphical method plots [A], ln[A] and 1/[A] against time and asks which is straight; the book calls the graphical or [[Least_squares|regression]] approach more reliable than a two-point estimate, and the variant sim under [[Rate_equation|Rate equation]] replots the peroxide and butadiene datasets under an assumed order until one picture becomes a line.[^af-graphical] ### Fast reactions Reactions complete in milliseconds or less cannot be started by mixing in a beaker. Stopped-flow instruments mix two solutions and record an optical signal as the mixture ages; flash photolysis starts a reaction with a pulse of light and watches the intermediates decay; relaxation methods disturb an equilibrium with a sudden jump in temperature or pressure and time the return. The 1967 Nobel Prize in Chemistry recognised the last two families of methods.[^nobel-1967] ## Equilibrium A reversible reaction A ⇌ P runs forward with rate constant kf and backward with kr, and at [[Dynamic_equilibrium|dynamic equilibrium]] the two rates are equal, so `K = kf/kr`.[^af-keq] For a reversible first-order reaction the approach to equilibrium is itself exponential: `ln(([A] − [A]eq)/([A]0 − [A]eq)) = −(kf + kr)·t`, so the distance from equilibrium halves every 0.693/(kf + kr), faster than either reaction alone.[^boyd-reversible] Starting from pure A, the material balance gives `[A]eq = [A]0/(1 + K)`. A catalyst that multiplies kf and kr by the same factor shortens the approach without moving [A]eq, the content of the Catalysis sim; a change in temperature moves both K and the rates, the content of [[Le_Chatelier's_principle|Le Chatelier's principle]]. ## Free energy The [[Gibbs_free_energy|Gibbs free energy]] change of a reaction decides its direction and the position of equilibrium, but says nothing about its speed; the Portal Book's thermodynamics chapter separates spontaneity from rate at the outset.[^af-spont] The classic illustration is diamond, whose conversion to graphite is spontaneous at room conditions and yet immeasurably slow, and a mixture of [[Hydrogen|hydrogen]] and [[Oxygen|oxygen]] is likewise far from equilibrium yet stable until a spark supplies the activation energy. Kinetics adds the barrier to the [[Thermodynamics|thermodynamic]] picture: the activation energy is the height of the transition state above the reactants, while the reaction [[Enthalpy|enthalpy]] is the height of the products above the reactants, and the two are independent.[^af-ea] A strongly exothermic reaction can be slow and a mildly endothermic one fast; only the barrier sets the rate. ## Applications and models Rate laws are the working equations of [[Chemical_engineering|chemical engineering]], atmospheric chemistry, [[Combustion|combustion]] science and biochemistry. A reactor is sized from the rate law and the required conversion; an ozone budget is a set of coupled rate equations for photochemical steps; a flame is a network of radical reactions; and an [[Enzyme|enzyme]] follows the saturating rate law of [[Michaelis–Menten_kinetics|Michaelis–Menten kinetics]], `v = Vmax·[S]/(KM + [S])`, this article's biochemistry door. In [[Electrochemistry|electrochemistry]] the same ideas appear as [[Electrochemical_kinetics|electrode kinetics]], where the electrode potential takes the place of temperature through the [[Butler–Volmer_equation|Butler–Volmer equation]]. Models of a reaction begin with its mechanism, the sequence of [[Elementary_reaction|elementary steps]] whose orders do follow their molecularity. The overall rate law is derived by identifying the [[Rate-determining_step|rate-determining step]] or by applying the steady-state approximation to short-lived intermediates; multistep mechanisms "more often than not" lack closed-form solutions, and the Portal Book presents the steady state as an empirical device to be checked against the exact solution.[^boyd-ssa] ### Numerical methods A mechanism is a system of coupled [[Ordinary_differential_equation|ordinary differential equations]], one per species, and beyond the simplest cases it is solved numerically. Explicit integrators such as the classical [[Runge–Kutta_methods|Runge–Kutta]] scheme step the concentrations forward in time, but a mechanism whose steps differ in speed by orders of magnitude is stiff: the step size is limited by the fastest process even when only the slowest is of interest, and implicit stiff solvers are used instead. The Wikitube steady-state sim is built that way: its two-step mechanism is integrated offline with a stiff solver and shipped as a table of trajectories, because at the fast end of its slider an explicit step would have to be smaller than about 2.8/(κ + λp) in scaled time to stay stable. Reaction networks that also move in space become [[Reaction–diffusion_system|reaction–diffusion systems]], and whole plants are modelled in [[Process_simulation|process simulation]] software that carries a rate law for every unit. ## See also - [[Rate_equation]] - [[Reaction_rate]] - [[Half-life]] - [[Arrhenius_equation]] - [[Catalysis]] - [[Steady_state_(chemistry)]] - [[Michaelis–Menten_kinetics]] - [[Chemical_equilibrium]] ## References [^af-ratelaw]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.3 Rate Laws, pp. 799, 804–805 (orders are experimental; k depends on temperature; units of k). https://open.umn.edu/opentextbooks/textbooks/chemistry-atoms-first [^af-halflives]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.4 Integrated Rate Laws, pp. 806–816 (integrated laws and half-lives for zero, first and second order; Table 17.2). [^boyd-integrated]: Boyd, W. Christopher (2025). *Exploring Inorganic and Organometallic Chemistry*. Chapter 10 Kinetics, pp. 275–277 (first- and second-order integrated laws; the second-order half-life depends on [A]₀). https://open.umn.edu/opentextbooks/textbooks/exploring-inorganic-and-organometallic-chemistry [^boyd-pseudo]: Boyd, W. Christopher (2025). *Exploring Inorganic and Organometallic Chemistry*. Chapter 10 Kinetics, pp. 279–280 (pseudo-order rate constants with one reactant in excess). [^boyd-reversible]: Boyd, W. Christopher (2025). *Exploring Inorganic and Organometallic Chemistry*. Chapter 10 Kinetics, pp. 275–276 (reversible first-order reaction A ⇌ P and its approach to equilibrium). [^boyd-ssa]: Boyd, W. Christopher (2025). *Exploring Inorganic and Organometallic Chemistry*. Chapter 10 Kinetics, pp. 280–282 (multistep mechanisms and the steady-state approximation). [^af-keq]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.7 Catalysis, pp. 834–835 (K = kf/kr; a catalyst does not change K). [^af-factors]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.2 Factors Affecting Reaction Rates, pp. 791–850 (page to pin). [^af-cyclobutane]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.4, pp. 806–807 (cyclobutane decomposition, k = 9.2 × 10⁻³ s⁻¹, 80.0 % decomposed; the printed time, about 175 s, is a derived value in the Wikitube extract). [^af-ea]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.5 Collision Theory and §17.6 Reaction Mechanisms, pp. 818, 828 (activation energy as the transition-state height; ΔH as the product–reactant difference). [^af-zero]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.4, pp. 812–813, 816 (zero-order decomposition of ammonia on tungsten; 0.0028 → 0.0001 M in 35 min; t½ = 18 min). [^af-zero-cond]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.4, p. 812 (zero order as conditional, "pseudo" behaviour). [^af-butadiene]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.4, pp. 809–810, 816 (butadiene, k = 5.76 × 10⁻² L mol⁻¹ min⁻¹; 0.200 M → 0.179 M in 10.0 min; the 86.8 min half-life at 0.200 M is derived from the book's k in the Wikitube extract). [^af-i131]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.4, pp. 807, 814 (iodine-131, k = 0.138 d⁻¹, t½ = 5.02 d). [^af-arrhenius]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.5 Collision Theory, pp. 819–820 (the Arrhenius equation and its linearized form; R = 8.314 J mol⁻¹ K⁻¹). [^af-rule]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.2 (the rule of thumb that rates roughly double per 10 °C; page to pin). [^af-hi]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.5, pp. 819–821 (hydrogen iodide rate constants 555–781 K; Ea = 1.8 × 10⁵ J/mol). [^af-catalyst]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.7 Catalysis, p. 828 (a catalyst lowers the activation energy of the rate-determining step by a different mechanism; reactant and product energies unchanged). [^af-paths]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.7, Example 17.15, pp. 828–829 (catalysed path: first-step barrier 70 kJ against 80 kJ; second step 20 kJ on both). [^af-h2o2]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.4, pp. 807–808 (hydrogen peroxide data at 6-hour intervals; the ln plot is linear, so first order). [^af-initial]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.3, pp. 801–804 (method of initial rates; NO + O₃ five trials, first order in each). [^af-graphical]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 17 Kinetics, §17.4–17.5, pp. 808, 816, 821 (the linearizing plots for each order; the graphical or regression approach is "more reliable"). [^af-spont]: Flowers, Paul; Neth, Edward; Robinson, William et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 12 Thermodynamics, §12.1 Spontaneity, pp. 597–622 (spontaneity is not speed; diamond to graphite; page to pin). [^nobel-1901]: The Nobel Prize. "The Nobel Prize in Chemistry 1901 — Jacobus H. van 't Hoff." NobelPrize.org, Nobel Prize Outreach. https://www.nobelprize.org/prizes/chemistry/1901/summary/ [^nobel-1903]: The Nobel Prize. "The Nobel Prize in Chemistry 1903 — Svante Arrhenius." NobelPrize.org, Nobel Prize Outreach. https://www.nobelprize.org/prizes/chemistry/1903/summary/ [^nobel-1909]: The Nobel Prize. "The Nobel Prize in Chemistry 1909 — Wilhelm Ostwald." NobelPrize.org, Nobel Prize Outreach. https://www.nobelprize.org/prizes/chemistry/1909/summary/ [^nobel-1967]: The Nobel Prize. "The Nobel Prize in Chemistry 1967 — Manfred Eigen, Ronald G. W. Norrish, George Porter." NobelPrize.org, Nobel Prize Outreach. https://www.nobelprize.org/prizes/chemistry/1967/summary/ ## External links - *Chemistry: Atoms First 2e* (OpenStax, 2019), Portal Book 051 — Open Textbook Library record: https://open.umn.edu/opentextbooks/textbooks/chemistry-atoms-first - *Exploring Inorganic and Organometallic Chemistry* (2025), Portal Book 052 — Open Textbook Library record: https://open.umn.edu/opentextbooks/textbooks/exploring-inorganic-and-organometallic-chemistry - The Wikipedia pair's External links section lists the pair's own links. <!-- MATTERSIM:BEGIN g33 — Matter & Energy Cluster microsim (framework build, specs/sims/Chemical_kinetics.json); do not hand-edit inside --> **Microsim — three.js (Wikitube framework):** *Chemical kinetics* <div class="wt-sim" data-src="https://wikitube-3d-microsims.netlify.app/matter/Chemical_kinetics.html" data-title="Chemical kinetics"></div> *Built from `MICROSIM_GUIDE/specs/sims/Chemical_kinetics.json`; part of the [[PORTAL_Matter|Matter portal]] spine (section sims and See-also variants).* <!-- MATTERSIM:END --> ## Wikipedia : Wikitube **Strict pair:** [Wikipedia](https://en.wikipedia.org/wiki/Chemical_kinetics) : [Wikitube](https://en.wikitube.io/wiki/Chemical_kinetics) · pinned revision [1351367206](https://en.wikipedia.org/w/index.php?oldid=1351367206) · 2026-09-11 ## Previous hub tags Hubs: `Life_Physics`. Portals: [[PORTAL_Chemistry]]. --- *Matter & Energy Cluster child articles, wave 1 · 2026-09-11 · drafted · Chemistry row K32 · sim pending (matter/Chemical_kinetics).*