AQA A Level Chemistry - Kinetics

• Particles are constantly moving in random directions • Consider this reaction (pictured) • Molecule A must collide with molecule B in order for them to react • Most of the time, their collisions are unsuccessful • This means they do not end up forming a product. Instead, they bounce off each other and go their separate ways • A successful collision means the particles react • Particles must have enough energy when they collide for the collisions to be successful

• The minimum energy needed for a reaction to occur/start

• Particles must collide • With E ≥ Ea • This symbol ≥ means greater than or equal to

• A small number of particles have E ≥ Ea

• A small number of particles have E ≥ Ea

• Molecules gain energy • Due to collisions

• Collisions cause some molecules to slow down or lose energy

• The rate of change in concentration • per unit of time • Units: mol dm-3 s-1

• Temperature • Concentration • Pressure (for gas based reactions) • Catalysts • Surface area (for solid reactants, as surface area increases, rate of reaction increases)

• More particles have E ≥ Ea • Increases frequency of successful collisions

• Increase in number of particles per unit volume • Increases frequency of successful collisions

• Increase in number of particles per unit volume • Increases frequency of successful collisions

• Lowers Ea • So there are more particles with E ≥ Ea • Increases frequency of successful collisions

• Increased number of reactant particles made available • Increases frequency of successful collisions

• Particles in a liquid/solution/gas have a range of kinetic energies • The range of energies of particles in a gas is represented by the Maxwell-Boltzmann distribution (see graph below) A - Line must start at the origin as no particles can have zero energy B - The peak of the curve shows the most probable value for the energy of the particles (Emp) C - The area under the curve shows the total number of particles D - The area under the curve, but to the right hand side of Ea represents the particles with enough energy to react E - The line levels out towards the x axis, but never touches it F - The mean energy of particles is a line where the area under the curve on either side of the line is equal • The distribution shows that only a small fraction of particles can collide with enough (E ≥ Ea) to start a reaction/break bonds (ie to allow the reactants to overcome the activation energy and form products)

• When drawing the curve for a higher temperature: • The High T peak must be further to the right • The curves must only cross once • The High T curve is wider/broader • The size of the shaded area is larger for High T peak • At high temperature/when temperature increases the rate of reaction increases because: • The proportion of particles with E ≥ Ea increases, so more frequent successful collisions • The rate of reaction approximately doubled when temperature rises by 10 degrees C

• A small increase in temperature leads to a significantly large number of particles with the activation energy

• When drawing the curve for a lower temperature: • The low T peak must be further to the left • The curves must only cross once • Low T curve is narrower • The size of the shaded area is smaller for Low T curve Increasing • At low temperature/when temperature decreases the rate of reaction decreases because: • The proportion of particles with E ≥ Ea decreases • So less frequent successful collisions

• An increase in concentration (or pressure in gases) increases the rate of reaction because: • Increase in number of particles per unit volume • So more frequent successful collisions • The rate of reaction doubles when concentration doubles: • The number of particles per unit volume doubles • So the number of frequent successful collisions also doubles

• The Emp and Ea would stay the same • The total area under the graph would increase

• A catalyst is a substances that increases the rate of reaction • but is not used up or changed • speeds up a reaction but is chemically unchanged at the end of the

• By providing an alternative reaction pathway • that has a lower activation energy

• The activation energy decreases • So more of the particles have a collision energy greater than the activation energy (E ≥ Ea) • So there are more frequent successful collisions