equilibrium + acids/bases Flashcards
chemical kinetics
the study of reaction rates
collision theory
for a reaction to occur, substance particles must physically collide in space
effective collisions
- a collision must be effective for a rxn to occur
- the molecules must collide in the correct orientation, and with sufficient energy
activation energy
the minimum energy required for a reaction to occur
transition state
- when bonds are broken, produces an unstable, temporary molecule
“activated complex” - as rxn continues, will settle into product
catalyst
a substance that reduces the activation energy of a reaction by changing reaction pathway without being consumed so that the transition state is different and lower in energy
rate of reaction
how fast a reaction progresses & proportional to the number of effective collisions per second
increase reaction rate by:
-more reactant surface area
-higher temperature
-a catalyst
-increasing concentration
surface area (rxn rate)
more surface area btw reactants=more collisions=more effective collisions=higher rxn rate
dissolving into solutions gets maxium surface area (liquids & gases)
temperature (rxn rate)
increase in temperature means an increase in kinetic E of the reactants=more reactants able to overcome the activtion E=higher rxn rate
concentration (rxn rate)
higher concentration of reactants=more collisions=more effective collisions=higher rxn rate
as a rxn progresses, reactants turn to products so concentration drops & rxn rate slows over time
pressure for gaseous systems (rxn rate)
increase in pressure=more particle-wall collisions=more particle-particle collisions=more effective collisions=higher rxn rate
as a rxn progresses, reactants turn to products so reactant pressure drops & rxn rate slows over time
catalyst (rxn rate)
adding a catalyst lowers the activation energy so more particles can overcome the activation barrier, and increases rxn rate
activation energy & rxn rates
higher activation energy (less collisions with enough energy to overcome the barrier) means slower rate & vice versa
reversible reactions
when the energies of the reactants & products are similar, no matter the amount of activation E, the rxn rates are similar & the reaction is reversible
irreversible reactions
if the energies of the reactants & products are different, the activation E of one will be a lot larger & the reaction will be irreversible
reaction rates in action
- forward reaction rates are higher when there’s more reactants
- reverse reaction rates are higher when there’s more product
- as one dwindles the other will become more of
- rates (not necessarily concentration of matter) will become equal at equilibrium
chemical equilibrium
when the rate of the forward reaction & the rate of the reverse direction are equal (not zero); reactants and products are still doing their reactions it’s just that there’s no net change
closed systems
to reach equilibrium, reactions must take place in a closed system where no matter can be exchanged with its surroundings (energy is fine)
Keq
a state of equlibrium: a ratio of products to reactants once a stable equilibrium is reached (@ a standard temperature)
Kc
a molarity Keq; pure solids & liquids not included, only for gaseous & aqeous reactions
the concentration of products raised to the power of their coefficients over the concentration of reactants raised to the power of their coefficients
Kp
only for gaseous reactions, a partial pressure Keq
the partial pressure of products raised to the power of their coefficients over the partial pressure of reactants raised to the power of their coefficients
temperature dependance
an equilibrium constant is specific for a certain temperature as the ratio of reactants and products change when temperature fluctuates
reaction quotient
Q is the ratio of reactant to product concentration at any other time but equlibrium (basically Kc but not at equlibrium)
Q<K
reaction goes to the right, and more product is created
