OZ - kinetics *02

the energy that pairs of molecules must possess to react when they collide

collision theory collide in the right direction (facing each other the right way) collide with the minimum amount of energy needed (activation energy)

reactants below products

reactants above products

shows how molecules in a gas dont all have the same amount of energy

curve starts at (0,0) cause no molecules have zero energy a few molecules are moving very slowly most molecules are moving at a moderate speed some molecules have more energy than the activation enthalpy —> these are the only ones that can react

particles, on average, have more kinetic energy and will move faster so a greater proportion of molecules will have the activation energy and be able to react • changes the shape of the curve —> it pushes to the right

particles on average have more kinetic energy and will move faster, so a greater proportion of molecules will have the activation energy and be able to react so the frequency of collision that result in a reaction will increase —> more successful collision per unit time (increases rate) also because more kinetic energy there are more collisions that take place (collide more are more likely to react when they collide)

increasing conc. of reactants in solution = the particles are closer together on average so they collide more often more collision = more chances to react

increase pressure of a gas = particles closer together on average they collide more often = more chances to react

catalysts lower the activation enthalpy by providing a different way for the bonds to be broken and remade activation enthalpy lower = more particles have enough energy to react

monitor loss of reactant or formation of product • vol. of gas produced • loss of mass • change in pH • temperature change • taking samples at regular intervals and analysing them by titration

use data from the experiment to plot a graph - time on x-axis, change on y-axis to find the reaction at a particular time • draw a tangent to the curve at that time: gradient of tangent = rate of reaction • rate units = change / unit of time

rate always fastest at the beginning as a reaction continues the conc. of reactants decreases so there will be less frequent collisions between reactant particles and the rate will decrease

to increase rate of reaction by providing an alternative reaction pathway with a lower activation enthaply catalyst chemically unchanged at the end of the reaction

homogenous heterogenous

a catalyst that is in the same state as the reactants

forms intermediates • speeds up rate by forming one or more intermediate compounds with the reactants. - –> the products are then formed from the intermediate compounds • activation energy needed to form intermediates and then products LOWER than activation energy needed to make products directly from reactants • the catalyst is reformed at the end and is chemically unchanged

profile would have two humps - lower than the uncatalysed reaction hump! intermediates are formed at the trough between the humps

line for activation energy moves back - so there are molecules that have enough energy to react in catalysed reaction but not enough to react when uncatalysed the curve stays the same

chlorine radical breakdown of ozone