Chemical reaction simulator for rate practicals

A chemical reaction simulator shows a reaction happening particle by particle, so you can see why changing the conditions changes how fast it goes. Our chemical reaction simulator, the Reaction Rate Simulator, runs five classic rate of reaction practicals as particles colliding in a box. Change the temperature, concentration, surface area or catalyst while it runs and watch the concentration chart get steeper or shallower, with the share of collisions that actually react measured as you go.

Try it

Reaction Rate Simulator

Heat, concentrate and catalyse particles to see collision theory

Open the simulatorGo to its page

What the simulator shows

• The particles: the real species named on the stage, such as H⁺ ions hitting a magnesium surface, with the products appearing as they react.
• An energy distribution: the Maxwell-Boltzmann distribution measured from the moving particles, with the activation energy marked and the share beyond it.
• Concentration against time: the reactant's concentration falling, with your last run dashed behind it so you can compare.
• Readings: the time, the concentration, the rate, the share of hits that reacted, the share predicted by e^−Ea/RT, and how rare a successful collision would be at the real activation energy.

The five practicals

A Custom reaction lets you set the activation energy yourself. The simulator gives the source for each figure, and is open that several are rough: metal and acid reactions are limited mostly by how fast the acid reaches the surface, so there is no agreed single value.

How to use the chemical reaction simulator

1. Open the Reaction Rate Simulator and pick a Practical under Reaction.
2. Under Conditions, set the temperature (0 to 100 °C for the practicals) and the concentration (0.2 to 2.0 mol/dm³), the surface area of any solid, and add the catalyst where the reaction has one.
3. Under Run, press play. Change a condition mid-run and the slope of the chart changes from that moment.
4. Reset and run again with one change: the last run stays dashed on the chart for comparison.

One thing to know: real activation energies are so high that, at the particles' speed, almost nothing would react while you watched. So the model divides each practical's barrier by 6 (thiosulfate's 47 kJ/mol becomes 7.8 kJ/mol), and shows the real figure alongside. The trends are right; the time scale isn't the bench's.

Why the rate changes: collision theory

A collision makes products only if it carries enough energy to get over the activation energy, Ea, and, for many reactions, if the particles meet the right way round. Anything that makes successful collisions happen more often speeds the reaction up. The share of collisions with enough energy is given by the Boltzmann factor:

Experiments to try

Experiment 1: temperature

Choose sodium thiosulfate at 20 °C. The predicted share of collisions with enough energy, at the model's barrier, is 4.0%. Raise it to 40 °C and it becomes 4.9%, and the particles move faster so they collide more often too. At the real barrier the readout shows the bigger change: from 1 in 2.4 × 10⁸ collisions at 20 °C to 1 in 7.0 × 10⁷ at 40 °C, about 3.4 times as many successful collisions. That is the main reason heating speeds reactions up so much.

Experiment 2: concentration

Keep the temperature fixed and double the concentration. There are twice as many particles in the same space, so they collide more often each second. The share of collisions that react stays the same; there are just more collisions, and the chart starts steeper.

Experiment 3: surface area

With magnesium, switch from ribbon to turnings to powder. The acid can only reach the outside of the metal, so breaking it into smaller pieces exposes more of it and the rate rises, though the total amount of hydrogen at the end is the same.

Experiment 4: a catalyst

Choose hydrogen peroxide and add manganese(IV) oxide. The catalyst gives a route with a lower activation energy: about 58 kJ/mol instead of 75 kJ/mol in the simulator's data. At the real barriers a successful collision goes from 1 in 2.4 × 10¹³ to 1 in 2.2 × 10¹⁰, about 1,000 times more often. Count the catalyst grains before and after: they aren't used up.

Experiment 5: the disappearing cross

Sodium thiosulfate and acid make a cloud of sulfur. The simulator times when there is enough sulfur to hide the cross under the flask, the Cross hidden at reading, just like the practical. Run it at two temperatures and compare: the hotter run hides the cross sooner, and 1 ÷ time is a measure of the rate.

Try reading energy distributions at two temperatures, then press Mark to check your answer in the Reaction Kinetics course.

What the simulator doesn't do

• Any reaction you type in. It runs its five practicals and one custom reaction, not a reaction of your choice.
• Reversible reactions and equilibrium. For Kc and rate equations, use the Rates and Equilibrium Calculator.
• Amounts and yields. For how much product forms, use the stoichiometry calculator.
• Real times. Because the barriers are scaled down, its seconds aren't bench seconds. Compare runs with each other, not with a stopwatch.

Common mistakes

• Saying particles collide "harder" at higher concentration. They collide more often; the energy depends on temperature.
• Saying a catalyst gives particles more energy. It provides a route with a lower activation energy.
• Thinking surface area changes the amount of product. It changes the rate, not the total.
• Forgetting the fraction with enough energy. Heating increases collision frequency a little but the share above Ea a lot.

Read next

Chemistry · Rates and equilibriumRate of reaction formula: mean rate, units and tangentsThe rate of reaction formula: mean rate from reactant used or product formed, units g/s, cm³/s and mol/s, rate from a tangent, the practical and examples.