module 3: equilibrium Flashcards
rate =
concentration/ time
rate of reaction calculated by
reactant us up/ product formed overtime
if gas produced increase of in volume, decrease in mass
if reactant acid increase in pH over time
collision theory
particles collide with enough energy to break existing bonds
collide in correct orientation
not all collisions are successful
higher rate = more successful collisions per second
increasing concentration
increases the rate of reaction
more particles per unit volume
more frequent collisions
rate of reactions from graphs
instantaneous -> tangent
catalyst
increases rate of reaction without being used up
lowers activation energy by providing an alternate pathway
“proceed via alternate route with lower activation energy”
homogeneous catalyst
same state as reactants
heterogeneous catalyst
different state from reactants
benefits of using catalyst
lower temperature and pressure
reduce energy demand -> less fossil fuels burnt -> less limited resources used up and CO2 produced
increases sustainability and reduces emissions
cost saving
Boltzmann distribution key notes
not symmetrical
does not cross x-axis (asymptote)
starts at origin, no particles have zero energy
total area= total number of gas particles
Boltzmann curve Ea
larger area= greater proportion of particles with an energy greater than Ea therefore faster rate of reaction
Boltzmann distribution temperature graph aided
higher temperature peak lower and shifted right
greater proportion of molecules can overcome activation energy
Boltzmann distribution temperature
as temperature increases
average energy of particles increases
peak distribution moves right
flatter
greater proportion of molecules can overcome activation energy
rate of reaction increases
Boltzmann distribution catalysts
distribution curve same
activation energy lower
greater proportion of molecules exceed new lower activation energy
on collision, more molecules react to form products
rate of reaction increases
Le Chateliers principle
system works to counteract change
dynamic equilibrium
closed system
rate of forward reaction= rate of reverse reaction
concentration of reactants and products remain constant
increasing temperature
position of equilibrium
endothermic reaction favoured
changing concentration
position of equilibrium
equilibrium shifts to oppose change
factors affecting rate of reaction
concentration
surface area
temperature
pressure
catalyst
pressure same no of moles on both sides
no change
increasing pressure
position of equilibrium
favours side with fewer moles of gas
heterogeneous catalysts notes
reactants absorbed onto surface
weak bonds formed between reactants and catalysts
bonds between reactant molecules weaken and break
new bonds form between reactant molecules
catalyst “poisoning”
impurities bind to the surface
e.g. lead and sulfur
regeneration
blowing hot air over a catalyst
oxidises chemicals on the surface
used to remove carbon after catalytic cracking
Kc =1
equilibrium position around halfway between reactants and products
Kc < 1
equilibrium favours reactants
Kc >1
equilibrium favours products
(high concentration of products to reactants)
homogeneous equilibrium
reactions in which reacting species are in the same phase
heterogeneous equilibrium
different phase
solids and liquids do not appear
equilibrium constant
[A] (concentration of reactants) decreases to a constant
[B] (concentration of products) increases from zero to a constant
no matter the starting composition of reactants and products, the same ration is achieved at equilibrium at constant temperature
catalyst equilibrium position
no effect-> forward and reverse both increased by same amount
equilibrium reached faster
Ea
minimum energy for reaction to occur
equilibrium law
for the reaction aA + bB ⇌ cC + dD
kc = ([C]c [D]d) / ([A]a [B]b)
kc is in terms of concentration
equilibrium law simplified
Kc= [products]/ [reactants]
when equation for reversible reaction given in opposite direction
equilibrium constant becomes reciprocal
