3. Rates of Reaction

• The concentration of the reactants in solution • The pressure of reacting gases • The temperature of the reaction • Surface area of solid reactants • The presence of a catalyst

• Add dilute hydrochloric acid to the conical flask • Use a delivery tube to connect this flask to an inverted measuring cylinder upside down in a water trough • Add calcium carbonate chips into the conical flask and close the bung • Measure the volume of gas produced in a fixed time using the measuring cylinder • Repeat with different sizes of calcium carbonate chips

• Measure 50 cm³ sodium thiosulfate solution into a flask. • Measure 5 cm³ dilute hydrochloric acid in a measuring cylinder. • Place a paper with a drawn cross under the flask. • Add the acid to the flask and start the stopwatch immediately. • Look from above and stop the stopwatch when the cross disappears. • Repeat with different concentrations of sodium thiosulfate by diluting with water.

• Dilute hydrochloric acid is heated to a set temperature using a water bath • Add the dilute hydrochloric acid into a conical flask • Add a strip of magnesium and start the stopwatch • Stop the time when the magnesium fully dissolves • Repeat at different temperatures and compare results

• Add hydrogen peroxide into a conical flask • Use a delivery tube to connect this flask to a measuring cylinder upside down in water trough • Add the catalyst manganese(IV) oxide into the conical flask and close the bung • Measure the volume of gas produced in a fixed time using the measuring cylinder • Repeat experiment without the catalyst of manganese(IV) oxide and compare results

• Has a steeper gradient at the start • Becomes horizontal sooner • Forms the same amount of product

• Increasing surface area → rate increases • Smaller pieces / powdered solid → reaction is faster

• Increasing concentration → rate increases • Decreasing concentration → rate decreases

• Increasing pressure → rate increases • Only applies to reactions involving gases

• Increasing temperature → rate increases • Lower temperature → reaction is slower

• Catalyst → increases the rate of reaction • Catalyst is not used up and is chemically unchanged at the end

• Increasing the concentration means that there are more reactant particles in a given volume • This causes more collisions per second • Leading to more frequent and successful collisions per second • Therefore, the rate of reaction increases

• Increasing the pressure means that there are the same number of reactant particles in a smaller volume • This causes more collisions per second • Leading to more frequent and successful collisions per second • Therefore, the rate of reaction increases

• Increasing the temperature means that the particles have more kinetic energy • This causes more collisions per second • Leading to more frequent and successful collisions per second • Therefore, the rate of reaction increases

• Increasing the surface area means that a greater surface area of particles will be exposed to the other reactant • This causes more collisions per second • Leading to more frequent and successful collisions per second • Therefore, the rate of reaction increases

A substance that increases the rate of a reaction, but is chemically unchanged at the end of the reaction

A catalyst works by providing an alternative pathway with lower activation energy.

The minimum amount of energy that reacting particles must have in order for a reaction to occur. The initial increase in energy, from the reactants to the peak of the curve, represents the activation energy.