topic 9: kinetics I Flashcards
what is the activation energy`
the minimum energy which particles need to collide to start a reaction
when can reactions only occur
when collisions take place between particles having sufficient energy. the energy is
usually needed to break the relevant bonds in one or either of the reactant molecules
what does the Maxwell Boltzmann energy distribution show
the spread of energies that molecules of a gas or liquid have at a particular temperature
what does the area under the curve represent
the total number of particles present
what does the Emp represent
the most probable energy (not the same as mean energy)
highest peak of the curve
at the beginning of the curve (steep part of the curve), what does it represent
A few have low energies because collisions cause some particles to slow down
why should the energy distribution go through the origin
because there are no molecules with no
energy
why should the energy distribution never meet the x axis
as there is no maximum energy for molecules
what do only a few particles have
energy greater
than the activation energy
what is the label of the y - axis
fraction of molecules with energy
how can a reaction go to completion if few particles have energy greater than EA
particles can gain energy through collisions
what happens as the temperature increases
the distribution shifts
towards having more molecules with higher energies
when temperature increases, why does the total area under the curve remain constant
because the total number of particles is constant
when temperature increases, what happens to the Maxwell Boltzmann distribution curve
it goes lower and to the right
what is the rate of reaction defined as
the change in concentration of a substance in unit time
- its usual unit is mol dm-3s-1
how do you work out the reaction rate
gradient of tangent to curve
in the experiment between sodium thiosulfate and hydrochloric acid, how is reaction rate normally measured
1/time where the time is the time taken for a cross placed underneath the reaction mixture to disappear due to the cloudiness of the sulfur
this is an approximation for rate of reaction as it does not include concentration. We can use this because
we can assume the amount of sulfur produced is fixed and constant
what happens when you increase concentration and pressure
there are more particles per unit volume and so the particles collide with a greater frequency
and there will be a higher frequency of effective collisions
what happens to the shape of the energy distribution curve when concentration increases
curves do not change (i.e. the peak is at the same energy) so the Emp and mean energy do not change
they curves will be higher, and the area under the curves will
be greater because there are more particles
more molecules have energy > EA
(although not a greater proportion)
what happens to the rate curves the higher the conc/ temp/ surface area
steeper gradient
what happens as the temperature increases
at higher temperatures the energy of the particles increases. they collide more frequently and more often with energy greater than the activation energy. more collisions result in a reaction
what is the effect of increasing surface area
will cause successful collisions to occur more frequently between the
reactant particles and this increases the rate of the reaction
what is a catalyst
a substance increases reaction rates without getting used up they do this by providing an alternative route or mechanism with a lower activation energy
what happens to the reaction rate when the activation energy is lowered
more particles will have energy > EA, so there will be a higher frequency of effective collisions. the reaction will be fast
what are heterogeneous catalyst
in a different phase from the reactants
what state are heterogenous catalyst normally in
usually solids
whereas the reactants are gaseous or in solution. the reaction occurs at the surface of the catalyst.
how does heterogeneous catalyst work
- adsorption of reactants at active sites on the surface
may lead to catalytic action - the active site is the place where the reactants adsorb on to the surface of the catalyst
- this can result in the bonds within the reactant molecules becoming weaker, or the molecules being held in a more reactive configuration
- there will also be a higher
concentration of reactants at the solid surface so leading to
a higher collision frequency
what is the effect of pressure on heterogeneous catalysts
- increasing pressure has limited effect on the rate of heterogenous catalysed reactions because the reaction
takes place on surface of the catalyst - the active sites on the catalyst surface are already saturated with
reactant molecules so increasing pressure wont have an effect
how are catalysts used industrially
speeds up the rate allowing lower temperature to be used
(and hence lower energy costs) but have no effect on equilibrium
what are the environmental benefits of catalysts
- catalysed reactions can occur at lower temperature so less fuel needed and fewer emissions from
fuels - catalysed reaction enables use of an alternative process with higher atom economy so meaning fewer raw materials needed and less waste products are produced
