A nail rusts for months, yet a mixture of hydrogen and oxygen explodes in a split second. Reaction speed comes down to molecular collisions: how often molecules meet and whether they have enough energy to get over the “hill” of activation energy. Heat the molecules, add a catalyst, feed an enzyme, race some beakers and shift the equilibrium in a syringe of brown gas.
Temperature, concentration, particle size and catalyst — four knobs a chemist turns to speed reactions up or slow them down.
Simplified model: a flat (2D) box, and an activation energy more than ten times smaller than in typical reactions — at a real 50–100 kJ/mol, a few hundred molecules would not react even once in hours. The molecules are also millions of times slower than in reality.
Van ’t Hoff's rule, “2–4 times for every 10 °C”, is only an approximation: it works for many reactions near room temperature when Eₐ is about 50–100 kJ/mol. At higher temperatures the same 10 °C step gives less and less.
Molecules have to collide. The collision has to be energetic enough to break the old bonds — the energy needed is the activation energy Ea. Often orientation matters too: the molecules have to hit each other in the right spots.
Higher concentration → more collisions per second, but the same fraction of effective ones. Higher temperature → faster molecules (slightly more collisions), but above all a larger fraction of collisions with energy above Ea — that's why the rate climbs so steeply.
At 100 °C higher (from 25 to 125 °C), molecules move on average only about 15% faster, yet a reaction with Ea = 50 kJ/mol speeds up about 160 times.
2 H₂O₂ → 2 H₂O + O₂, ΔH ≈ −98 kJ per mol of H₂O₂. Choose a catalyst:
The balls get energy from random collisions. To make anything visible, their energy in the animation is about 13 times the average molecular energy at that temperature; the “Speed-up” is calculated for the real temperature.
Inside the exhaust pipe sits a ceramic “honeycomb” with thousands of tiny channels coated with a thin layer of platinum, palladium and rhodium. On that surface CO is oxidized to CO₂, unburned hydrocarbons to CO₂ and H₂O, and nitrogen oxides NOₓ are reduced to N₂. It only works once it's warm (roughly above 250–300 °C), which is why a car releases the most pollution in the first few minutes of a drive.
Hydrogen peroxide decomposes on its own very slowly. Yeast (the enzyme catalase) or potassium iodide makes oxygen come out within seconds and foam up the dish soap. This is a teacher demo for a school lab: concentrated H₂O₂ burns skin and eyes, and the mixture gets very hot.
Ammonia NH₃ is made from nitrogen in the air and hydrogen on an iron catalyst, at about 150–300 bar and about 400–500 °C. Without the catalyst, the N≡N triple bond would break far too slowly.
Ammonia is turned into nitrogen fertilizers — by some estimates they make it possible to feed roughly half of humanity. In Poland, ammonia is made this way at the nitrogen plants in Puławy and Tarnów, among others.
A catalyst is not used up and changes neither ΔH nor the position of equilibrium — it speeds up the reaction in both directions, so equilibrium is simply reached sooner.
Michaelis–Menten curve: at high [S] all the enzymes are busy and the rate approaches Vmax. KM is the concentration at which v = Vmax/2.
An enzyme is a protein with a pocket — the active site. Only a substrate with the right shape and charge distribution fits into it (the “lock and key” model). Today we know the active site adjusts slightly as the substrate enters — this is induced fit.
A competitive inhibitor imitates the substrate and occupies the active site. Lots of substrate can “push it out” — Vmax stays the same, only KM grows.
At high temperature the protein chain unravels (denaturation) and the active site loses its shape. In the model the enzyme recovers when cooled; real proteins usually don't — a cooked egg never turns runny again.
“Biological” detergents contain enzymes: proteases (protein — blood, grass), amylases (starch) and lipases (fat). They work best at about 30–40 °C; in boiling water they would denature.
Fresh pineapple (bromelain) and kiwi (actinidin) contain enzymes that cut proteins — gelatin too, so jelly made with them won't set. Canned pineapple works: heating denatured the enzyme.
Once the apple is cut, the enzyme polyphenol oxidase oxidizes phenolic compounds with oxygen from the air, producing brown pigments. Lemon juice slows this down: it lowers the pH and acts as an antioxidant (vitamin C).
In the fridge (about 4 °C), spoilage reactions and bacterial growth are several times slower than at 20 °C. In the freezer (−18 °C) microbes don't multiply and the water is ice — but food still changes slowly: fats go rancid and meat “dries out” in the frost.
Chew a piece of bread for a while: amylase in saliva cuts starch into shorter sugars, including maltose, and after a moment it tastes slightly sweet. In the stomach (pH about 2) amylase stops working and pepsin takes over.
Many drugs are enzyme inhibitors. Aspirin permanently blocks cyclooxygenase (an enzyme that makes compounds involved in inflammation and blood clotting), and some blood pressure drugs inhibit angiotensin-converting enzyme.
In each flask: 0.40 g of marble (CaCO₃, 4 mmol) and 50 mL of hydrochloric acid — the acid is always in excess, so every flask eventually gives about 96 mL of CO₂. Only the rate changes. Simplified model: rate ∝ surface area × acid concentration × Arrhenius factor.
| Flask | mL/s | 50 mL | CO₂ |
|---|
In a first-order reaction the rate is proportional to concentration. Then half of the substance always disappears in the same time — the half-life — no matter how much there was. Radioactive decay behaves the same way.
In a zero-order reaction a fixed amount disappears every hour. That's roughly how the liver removes alcohol: the enzymes are “jammed”, so they run at full speed — the usual figure is about 0.1–0.2 ‰ (per mille) per hour, and it varies from person to person.
A log burns for hours, but sawdust, flour, powdered sugar or coal dust dispersed in the air can burn up in a split second. The huge surface area of the tiny particles is in contact with oxygen — this is a dust explosion, a danger in mills, silos and mines.
Granulated sugar dissolves in tea faster than a sugar cube, and powdered sugar faster still. Same mass, more surface in contact with the water. In the same way, finely split kindling gets a campfire going, while thick logs keep it burning for a long time.
Chewing breaks food into small pieces and mixes it with saliva: digestive enzymes get more surface to work on. That's why chewing thoroughly helps digestion, and big bites are digested more slowly.
Forming N₂O₄ is exothermic (ΔH ≈ −57 kJ/mol). Real equilibrium sets in within a fraction of a second — here it's slowed down. We ignore the condensation of N₂O₄ (it boils at about 21 °C at 1 atm). NO₂ is toxic: this experiment is shown in sealed glass ampoules.
Ammonia synthesis is exothermic and reduces the number of moles of gas (4 → 2). So Le Chatelier's principle says: high pressure and low temperature shift the equilibrium toward NH₃.
But at low temperature the reaction — even with an iron catalyst — runs so slowly that the plant would produce too little. So a compromise is chosen: about 400–500 °C and about 150–300 bar, and the unreacted gas is cycled back into the reactor.
Cobalt(II) chloride in weather “barometers” and humidity indicators: blue when it's dry, pink when it's humid — because water shifts the equilibrium between forms of the complex.
Graph: ideal gas, equilibrium constant from an empirical formula; numbers are approximate. Rate — a qualitative model.
To rust, iron needs oxygen and water at the same time; salt speeds things up because its ions help current flow in tiny cells on the metal's surface. The pace in the model is arbitrary, but the order is the same as in the classic school experiment.
The reaction rate tells you how fast reactants are used up or products appear — e.g. in mol/(L·s) or in mL of gas per second. It can be measured in many ways: by gas volume, or by changes in mass, color or conductivity.
Only molecules that collide with an energy of at least Ea react (and usually only in the right orientation). At the top of the “hill” an unstable activated complex forms — a transition state in which the old bonds are already breaking and the new ones are only forming. The fraction of molecules with energy above Ea rises very steeply with temperature:
For Ea = 50 kJ/mol, warming from 25 to 35 °C speeds the reaction up about 1.9 times. Hence the school “van ’t Hoff rule”: 2–4 times for every 10 °C — only an approximation.
A catalyst opens a different reaction pathway with a lower activation energy. It isn't used up and changes neither ΔH nor the equilibrium constant — it speeds up the forward and reverse reactions equally. Enzymes are biological catalysts (usually proteins): extremely fast and selective, but sensitive to temperature and pH.
When a solid reacts with a liquid or gas, only the molecules on the surface react — grinding it up multiplies the surface area many times. Higher concentration means more collisions per second. In a first-order reaction the concentration falls exponentially and has a constant half-life:
In a reversible reaction reactants form products, and products turn back into reactants. When the two rates become equal, the concentrations stop changing even though the molecules keep reacting — this is dynamic equilibrium. For A ⇌ B with rates k₁[A] and k₂[B]:
Le Chatelier's principle: a system at equilibrium that is disturbed shifts so as to partly counteract the change. Heating favors the endothermic direction, compression favors the side with fewer moles of gas, and adding a reactant favors forming products.
Simplifications: the simulations are two-dimensional and slowed down, the energies in the molecule box are rescaled, and the beaker race and rust are simple models with approximate numbers.