Reaction rates (3.2.2)

Collision theory states that for a chemical reaction to take place, particles must collide in the correct orientation and with sufficient energy.

Activation energy, is the energy barrier that must be overcome for the reaction to proceed. It is typically measured as the energy difference between the reactants and the transition state.

This means that for a successful collision to occur, the reactant particles must collide with energy equal to or above the activation energy.

If reactant particles collide with sufficient energy (equal to at least the activation energy, ) and in the correct orientation they will react, leading to an effective collision.

If reactant particles collide with insufficient energy, they bounce off of each other and it will be an ineffective collision.

If reactant particles collide in the incorrect orientation, they will bounce off each other resulting in an ineffective collision.

If the concentration of solutions is increased there are more particles per unit volume.

This will increase the frequency of collisions, and therefore the frequency of effective collisions will also increase.

In the diagram above, the concentrations of the reactant particles have increased. This would increase the frequency of effective collisions , and therefore increase the rate of reaction.

If the pressure of gaseous reactants is increased it has a similar effect to increasing the concentration.

Although the number of particles remains the same, the volume is decreased, and therefore the gaseous particles are more tightly packed. There are more gaseous particles per unit volume.

Increasing the number of particles per unit volume will increase the frequency of effective collisions, leading to an increase in the rate of reaction between the gaseous reactants.

Catalysts increase the rate of a chemical reaction but are not used up by the overall reaction.

Catalysts increase the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy.

This can be shown on enthalpy profiles.

Although catalysts increase the rate of a chemical reaction they do not impact the total frequency of collisions. However, they do increase the proportion of successful collisions and therefore the frequency of successful collisions.

Reactions can be monitored by measuring the mass change over time.

Procedure:

• Weigh the reaction vessel empty.
• Add the reactants and record the initial mass.
• At specific time intervals, reweigh and record changes.

Mass loss indicates consumption of solid or liquid reactants and formation of a gaseous product.

Mass gain implies consumption of a gaseous reactant and incorporation in a solid or liquid product.

In the reaction of solid magnesium with aqueous hydrochloric acid, hydrogen gas is produced, leading to a decrease in the mass of the reaction vessel.

Changes of gas volume during a reaction reflect consumption of gaseous reactants, or formation of gaseous products.

Use a gas syringe or displacement method to measure gas volume changes.

Record initial volume and measure at specific time intervals.

Hydrochloric acid reacts with magnesium to produce hydrogen gas, leading to an increase in gas volume as the reaction proceeds. Gas collection can be used to monitor the reaction rate.

Time is a crucial parameter for monitoring reaction rates.

Use a stopwatch or digital timer to record the time taken to reach specific reaction milestones. Accurate timing ensures precise determination of reaction rates.

The time taken for the milestone to be reached is inversely proportional to the reaction rate. This means that longer times indicate slower rates.

The milestone used will depend on the reaction but could include a colour change, onset of gas evolution, or a predefined temperature change.

In the reaction between sodium thiosulfate and hydrochloric acid, a precipitate of sulfur is formed, causing the solution to become cloudy. The reaction milestone is the point at which a covered cross is obscured.

The graph below illustrates how to draw a tangent and extract the relevant data from a concentration–time graph in order to calculate reaction rate.

The x-axis for time is often in minutes. When calculating the it is important to convert the time to seconds.

The gradient will commonly have the units of .

This method can also be used if there is a change of mass, change in volume etc. If this is the case, the units for the gradient will be different.

The Maxwell-Boltzmann distribution curve is a graphical representation showing the distribution of energies of molecules at a particular temperature.

The Maxwell-Boltzmann distribution curve plots the number of molecules on the axis and energy of the molecules on the axis.

The area under the curve in the Boltzmann distribution is equal to the number of molecules in the substance.

Upon changing conditions, the total area under the curve does not change.

The peak of the curve represents the most probable energy, ; the mode.

The curve is asymptotic (does not reach the axis), as molecules have no maximum kinetic energy, and starts at the origin , as molecules must have a non-zero energy.

The curve becomes useful in explaining reaction kinetics, as only molecules with energy equal to or greater than the activation energy will result in effective collisions. This is the area under the curve to the right of the line.

Increasing the temperature shifts the Maxwell-Boltzmann distribution curve to the right and down.

The total area under the curve remains unchanged, meaning a shift to the right must be accompanied by a lower peak height.

For decreasing temperature, the curve shifts to the left and the peak height increases.

The Maxwell-Boltzmann distribution shows how a small increase in temperature, from to , results in a larger proportion of molecules having sufficient energy to overcome the activation energy barrier and react.

Since more molecules are able to collide with sufficient energy to, the frequency of successful collisions increases, leading to a faster reaction rate at higher temperatures.

The use of a catalyst has no effect on the energy distribution or the shape of the Maxwell-Boltzmann distribution curve.

Catalysts lower activation energy so the position of the activation energy is shifted to the left.

This means that a greater proportion of molecules have an energy exceeding the activation energy and are able to react; a greater number of collisions per unit time are effective, and there is a faster rate of reaction.