Kinetics Flashcards
chemical kinetics
speed or rate of a reaction and its mechanism (do not confuse with the equilibria or extent)
rate
the change in concentration of reactants or products overtime (mol/L/s)
rate equation
rate = - 1/a (delta [A]/delta t) = -1/b (delta [B]/delta t) = 1/c (delta [C]/delta t) = 1/d delta [D]/delta t)
concentration and rate
- substances must come into contact to react
- more particles collisions, more reactions
- rate usually increases with concentration
physical state and rate
- for heterogeneous reactions (two different phases) collisions between reactants only occur at interfaces between phases
- number of collision between the reactants per unit time depends on surface area of more condensed phase
- increases surface area!!
temperature and rate
- increase temperature, increases average kinetic energy of particles
- as KE increases, particles move faster and collide more frequently with greater energy
- increases rate
catalysts and rate
catalysts participate in a chemical reaction and increase the rate of the reaction withoutt undergoing a net chemical change itself
- highly selective
- often determine product by only speeding up one aspect of a reaction
average rate
the reaction rate between two points
instantaneous rate
using a tangent you can find the rate or gradient at any point
rate law
rate = - 1/a (delta [A]/delta t) = k [A]^n
rate: mol/L/sec
k is the rate constant and is characteristic of conditions
n is the rate order and is determined experimentally
first order
rate directly proportional to concentration of a reactant
k[A]^1
- if k is doubled, [A] is doubled
- if [A] decreases, rate decreases
- plotted against time, it is non-linear
- rate is fastest at beginning
- k is s-1
second order
rate is proportional to the square of [A]
k[A]^2
if [A] doubles, rate quadruples
initially faster than first and then slows
k is L/mol/s
zero order
rate is independent of concentration
k[A]^0 = k
- [reactant] graph is linear of -k
- [product] is linear of +k
molL/s
overall reaction order with more than one reactant
k[A]^n[B]^m
overall = n+m
determining rate order: method of initial rates
- preform the reaction a number of times and vary the conditions (iodine clock)
- look at if the rate doubles, quadruples etc. when concentration is changed
- can determine k with algebra
determining rate order: integrated rate law
use if:
- initial concentration of a reactant
- concentration of a reactant overtime
- several measures of [] between times
integrated rate law: zero order
[A]t = [A]0 -kt
linear [A] vs time
integrated rate law: first order
ln[A]t = ln[A]0 -kt
area under graph decreases linearly
linear if you plot ln[A] against t
integrated rate law: second order
1/[A]t = 1/[A]0 + kt
linear if 1/[A] vs t
half life for first order
half life: time it takes to halve a concentration to halve
t1/2 = ln2/k
- independent of [A]
- successive half lives do not change as [A] changes
radioactive decay
the emission of a particle or photon that results from the spontaneous decomposition of the unstable nucleus of an atom
- loss of a particle
- loss of beta particle
- emission of y radiation
- rate is independent of chemical and physical form of the isotopes or temperatures
- first order process
- isotopes with shot half lives decay faster
carbon dating
- living things have constant C14:C12 ratio
- C14 becomes N14 + beta- when they die
- the half life is 5700 +/- 30 years
- comparing ratio with that of living organisms helps determine age
half life for second order
t1/2 = 1/k[A]0
half life for zero order
t1/2 = [A]0/2k
Arrhenius equation
k = Ae ^-(Ea/RT)
k is the rate constant A is the frequency factor Ea is activation energy R is 8.314J/k/mol T is temp in k
ln(k2/k1) = -Ea/R(1/T2 - 1/T1)
activation energy
the minimum energy required for a collision between molecules to result in a chemical reaction
- energy to overcome electrostatic repulsion and a minimum amount of energy to break bonds so new ones can form
- determines the rate of reaction
- molecules that collide with less than Ea bounce off each other and are chemically uncharged (only direction and speed altered)
- molecules with correct orientation (symmetry) react
Boltzman distribution of KE
- shaded areas (curve to baseline) are proportional to the total fractions of coliisons that involve the min Ea or more
- at higher T, more molecules are present with sufficient KE for reactions - increases rate
- area under curve is same for specified number of molecules
- Ea is not temp dependent
net reaction rate
sum of forwards and back reactions
rate = k1[A] - k-1[B]
catalysis
alters the rate of a reaction without appearing in any of the products by providing a new energy pathway which may have different Ea
homogenous catalysts
organometallic compounds and enzymes
- catalyst is in same phase as reactants
- collisions with reactants at ,ax because of uniform dispersion
- no common mechanism
heterogeneous catalysts
- Haber process uses alpha-Fe catalyst
- Zeolites for cracking petroleum
- at least one of the reactants with the solid surface (adsorption) so that chemical bond in the reactant becomes weaker and breaks
autocatalysis
one product catalyses the reaction eg. oxygen binds to haemoglobin
homogenous
enzymes react in aqueous solutions within a cellular compartment
heterogeneous
catalysts are embedded in the membrane
enzymes are
specific
- at low substrate concentrations, rate is first order with respect to [S]
- at high concentrations rate is zero order with respect to [S]
enzyme graph
steep increase and then flattens
- V max is the fast, flat area
- Km is half of Vmax
rate = V = (Vmax [S]/Km + [S])
reaction mechanisms and elementary reactions
reactants to products (may be many steps)
- can predict mechanism with elementary reactions
- test prediction with experimentally determined kinetics data
intermediates
- produced in one step and used in another
- do not appear in balanced equations
- help determine rate
rate determining steps
- slowest elementary steps
- orders of the rate equation are those from the rate determining step
- have the highest Ea barrier
molecularity
the number of molecules or atoms of reactant taking place
- unimolecular, bimolecular, trimolecular (rare) - harder to orientate
unimolecular (A to products)
first order
2A to products
bimolecular and second order
A + B
bimolecular and second order
2A + B
A + B + C
trimolecular and third
