3.1.5 Kinetics

Collision theory

Chemical reactions occur when particles collide with the correct orientations and with specific energy which is why many collisions do not lead to a reaction

The activation energy is the minimum amount of energy needed for a reaction to occur- the energy is used to break existing bonds and to initiate a reaction

Factors affecting rate of reaction: increasing surface area

• More particles are exposed so more particles are available to collide and react

• This leads to a higher frequency of successful collisions between the solid and other reactants

• More particles react in a given time, increasing the reaction rate

Factors affecting rate of reaction: increasing concentration

• More reactant particles in a given volume

• Particles on average will be closer together

• More collisions with correct orientation and sufficient energy

• This increases the frequency of successful collisions

• So faster reaction rate

Factors affecting the rate of reaction: pressure

• The gas particles are closer together so more particles in a given volume

• Therefore they will collide more frequently, increasing the rate of successful collisions

• Therefore increasing rate

Factors affecting the rate of reaction: increasing temperature

Increases the kinetic energy and the speed of the particles

As the particles move faster they collide with more frequently, increasing the rate of successful collisions

Kinetic energy also means that more particles have the necessary energy needed to overcome the activation energy, resulting in greater proportion of successful collisions

More collisions per unit time, so increased rate

Factors affecting the rate of reaction: presence of a catalyst

Provides an alternative pathway with a lower activation energy

Particles require less kinetic energy to react

More particles have energy greater than the activation energy and so can react

But the position of equilibrium does not change because it increases the rate of the forward and reverse reactions equally

Does not get used up in a reaction

The Maxwell Boltzmann distribution

Shows the variation in the kinetic energy of gas molecules in a reaction

Shows the spread of energies that molecules of gas have at specific temperatures

The number of particles is on the y axis, and kinetic energy on the x axis

Features of Maxwell Boltzmann distribution

1. The curve always goes through the origin because there are no molecules with zero energy

2. As the gradient rises: few have low energies because collisions cause some particles to slow down

3. The most probable energy is the most common kinetic energy level among molecules before gradually decreasing

4. At the average or mean energy, is NOT the peak of the curve as the energy of all the molecules is higher than the peak energy

5. The total area under the curve is equivalent to the total number of molecules

6. Only a few particles have energy greater than the activation energy

7. The curve never meets the x axis, as there is no maximum energy for molecules/ no upper limit for how fast particles can move

Effect of temperature on the distribution

At a lower temperature the curve shifts to the left

As a higher temperature, the curve shifts to the right

As temperature increases a greater number of particles have energy greater than the activation energy

The areas remain constant, as the number of molecules is constant, meaning that the peak height is lower

Affect of pressure on the distribution

In pressure does no change the kinetic energy of the particles, so the most probable energy and the activation energy stay the same

If pressure increases, the curve gets taller because there are more molecules of gas in a given volume, so the total area under the curve increases

So if pressure decreases, the curve becomes shorter

Affect of a catalyst on distribution

The regional high activation every appears as a vertical line of the far right

A catalyst moved this activation energy line further to the left

This means that more molecules have sufficient energy to react

This results in a faster rate of reaction

Types of catalysts

Homogenous catalysis

• The catalyst is in the same phase as the reactants (like dissolved ions in the same solution)

• Works by forming temporary intermediate compounds that break down more easily

Heterogenous catalysis:

• The catalyst is in a different phase from the reactants (like as solids metal surface with gas flowing over it)

It works in three steps:

1. Reactants stick to the surface (adsorption)

2. Reaction happens on the surface

3. Products leave the surface (desportion)

Required practical three: how the rate of reaction changes with temperature