Ch. 5: Chemical Kinetics Flashcards
what does change in Gibbs free energy (delta G) denote?
determines whether or not a reaction will occur by itself without outside assistance (spontaneous or not)
defn: reaction mechanism
proposed pathways/ the series of steps that proceed in order for a reaction to occur
sum of the steps gives the overall reaction
defn: rate-determining step
the slowest step in any proposed mechanism
acts like a kinetic bottleneck, preventing the overall reaction from proceeding any faster than that slowest step
defn + char + implications: collision theory of chemical kinetics
the rate of a reaction is proportional to the number of collisions per second between the reacting molecules
- not all collisions result in a chemical reaction
- an effective collision (one that leads to product formation) only occurs if the molecules collide with each other in the correct orientation and with sufficient energy to break their existing bonds and form new ones
defn + aka: activation energy (Ea)
part of the collision theory of chemical kinetics
the minimum energy of collision necessary for a reaction to take place
aka: energy barrier
eqn: rate of a rxn (collision theory of chemical kinetics)
rate = Z x f
Z = total number of collisions occurring per second
f = the fraction of collisions that are effective
eqn + defn: Arrhenius equation (collision theory of chemical kinetics) + defn: frequency factor
a more quantitatively rigorous analysis of the collision theory
k = Ae^ (-Ea/RT)
k = rate constant of a reaction
A = the frequency factor
Ea = activation energy of the reaction
R = ideal gas constant
T = temperature (Kelvin)
defn: frequency factor = the attempt frequency of the reaction = a measure of how often molecules in a certain reaction collide (unit = s ^-1)
what relationships can you infer from the Arrhenius equation? (k = Ae^(-Ea/RT)
As A increases, k increases
reaction rate increases with temperature because exponent becomes smaller (becomes less negative)
how can the frequency factor (A) be increased?
by increasing the number of molecules in a vessel
more molecules = more opportunities for collision
expl: transition state theory
when molecules collide with energy equal to or greater than the activation energy, they form a transition state in which the old bonds are weakened and the new bonds begin to form
the transition state then dissociates into products, fully forming the new bonds
defn: reaction coordinate (transition state theory)
traces the reaction from reactants to products, through the transition state
defn + aka + char: transition state
aka: activated complex
a theoretical intermediate between reactants and products
has greater energy than both the reactants and the products
is denoted by sort of a double cross symbol
once it is formed it can either dissociate into products or revert to reactants without any additional energy input
what is the difference between transition states and reaction intermediates?
transition states are theoretical constructs that exist at the point of max energy, rather than reaction intermediates which are distinct identities with finite timelines
defn: free energy diagram
illustrates the relationship between the activation energy, the free energy of the reaction, and the free energy of the system
defn: free energy of the reaction (deltaGrxn)
the difference between the free energy of the products and the free energy of the reactants
defn: exergonic and endergonic reaction (in terms of free energy)
EXERGONIC REACTION = a negative free energy change // energy is given off
ENDERGONIC REACTION = a positive free energy change // energy is absorbed
where is the transition state in the energy diagram?
at the peak of the energy diagram
what are the four factors affecting reaction rate in chemical kinetics?
- reaction concentrations
- temperature
- medium
- catalysts
how does reaction concentration affect reaction rate?
the greater the reactants’ concentrations, the greater the number of effective collisions per unit time
this leads to an increase in frequency factor (A)
SO reaction rate will increase for all but zero-order reactions
how does temperature affect reaction rate?
for nearly all reactions, the reaction rate will increase as the temperature increases
substance temperature = a measure of the particles’ average kinetic energy
increasing the temperature increases the average kinetic energy of the molecules
also the proportion of reactants gaining enough energy to surpass Ea (and thus can undergo reaction) increases with higher temperature
ALL REACTIONS ARE TEMPERATURE DEPENDENT AND EXPERIENCE AN OPTIMAL TEMPERATURE FOR ACTIVITY
what temperature is an enzymatic reaction optimal at? what happens after this optimal point?
35 - 40 deg. C (body temperature)
after 40 deg, the curve falls sharply because denaturation has occurred
are polar or nonpolar solvents generally preferred in terms of reaction rate?
polar –> bc the molecular dipole tends to polarize the bonds of the reactants, thereby lengthening and weakening them, allowing the rxn to occur faster
defn + char: catalyst
defn: substances that increase reaction rate without themselves being consumed in the reaction
char:
- interact with the reactants by adsorption or through the formation of intermediates
- stabilize reactants so as to reduce the activation energy necessary for the reaction to proceed
- return to their original chemical state upon product formation
how can catalysts affect reaction rate?
- increase the frequency of collisions between the reactants
- change the relative orientation of the reactants (making higher % of the collisions effective)
- donate electron density to the reactants
- reduce intramolecular bonding within reactant molecules
defn: homogeneous vs. heterogeneous catalysis
HOMOGENEOUS = the catalyst is in the same phase (solid, liquid, gas) as the reactants
HETEROGENEOUS = the catalyst is in a distinct phase as the reactants
what impact does a catalyst have on a reaction? what impact DOESN’T a catalyst have on a reaction?
DOES = decrease energy of activation for forward and reverse reactions (change rates of reaction (reverse and forward by the same factor) = MAKE SPONTANEOUS REACTIONS MORE MORE QUICKLY TOWARD EQUILIBRIUM
DOESN’T = no impact on free energies of reactants, products, or difference between them (no impact on equilibrium position or Keq) = DO NOT TRANSFORM NONSPON INTO SPON RXNS
defn + eqn + units: rate for the general reaction aA + bB –> cC + dD
rate = -delta[A]/adeltat = - delta[B]/bdeltat = delta[C]/cdeltat = deltaD/ddeltat
mol/L*s or M/s
we can describe the rate in terms of either the disappearance of reactants over time or the appearance of products over time
what is the first thing to look for whenever a question asks to determine the rate law for a reaction?
experimental data!
usually a chart that includes the initial concentrations of the reactants and the initial rates of product formation as a function of the reactant concentrations (often for 3-4 trials)
eqn + unit: rate law for the general reaction aA + bB –> cC + dD
rate=k[A]^x[B]^y
for nearly all forward, irreversible reaction
k = reaction rate coefficient/rate constant/proportionality constant
x, y = orders of the reaction, experimentally determined
units: concentration over time = M/s
the exponents x and y can be used to state the order of the reaction with respect to each reactant or overall: x is the order with respect to reactant A, y is the order with respect to reactant B, the overall order of the reaction is the sum of x and y
4 common traps to pay attention to in chemical kinetics
- assumption: x and y are the same as the stoichiometric coefficients
TRUTH: On the MCAT, the values of x and y are almost never the same as the stoichiometric coefficient
- there are only 2 cases in which they match
1. reaction mechanism is a single step and the balanced overall reaction is reflective of the entire chemical process
2. the complete reaction mechanism is given and the rate-determining step is indicated (although this can get complicated)
- mistaking the equilibrium constant for the rate law
TRUTH:
- equilibrium expression includes the concentrations of all the species in the reaction, both reactants and products
- expression for chemical kinetics includes only the reactants
- Keq indicates where the reaction’s equilibrium position lies
- the rate indicates how quickly the reaction will get there
- TRUTH: technically k is not a constant bc its particular value for any specific chemical reaction will depend on the activation energy for that reaction and the temperature at which the reaction takes place
for a specific reaction, at a specific temperature, the rate constant is indeed a constant
- assumption: the notion and principles of equilibrium apply to the system only at the end of the reaction
BUT the reaction rate can theoretically be measured at any time, but is usually measured at or near the beginning of the reaction to minimize the effects of the reverse reaction
how to use provided data in the experimental determination of rate law?
- identify a pair of trials in which the concentration of one of the reactants is changed while the concentrations of all other reactants remain constant
- with this, any change in the rate of product formation from one trial to another is fully attributable to the change in concentration of that one reactant
- repeat this process for the other reactant, using data from a different pair of trials, making sure that the concentration of only the reactant we are trying to analyze is changed from one trial to the other while the concentrations of all other reactants remain the same
- once the orders of the reaction have been determined with respect to each reactant, we can write the complete rate law, replacing the exponents x and y with actual numbers
- to determine the value of k, plug in actual values from any one of the trials (whichever is most convenient)
see ex. on pg. 179
defn: zero-order reaction
one in which the rate of formation of product C is independent of changes in concentrations of any of the reactants, A and B
constant reaction rate equal to k
rate = k[A]^0[B]^0 = k with units M/s
what are the two ways to change the rate of a zero-order reaction?
- temperature
- the rate constant itself is dependent on temperature - addition of a catalyst
- lowers the activation energy, thereby increasing the value of k
defn + implication: first-order reaction
has a rate that is directly proportional to only one reactant, such that doubling the concentration of that reactant results in a doubling of the rate of formation of this product
rate = k[A]^1 or rate = k[B]^1
where k has units s^-1
a first order rate law with a single reactant suggests that the reaction begins when the molecule undergoes a chemical change all by itself, without a chemical interaction, and usually without a physical interaction with any other molecule
what is a classic example of a first-order reaction? + eqn
radioactive decay
[A]t=[A]0e^(-kt) = the concentration of radioactive substance A at any time t
defn + implications: second-order reaction
has a rate that is proportional to either the concentrations of two reactants or to the square of the concentration of a single reactant
rate = k[A]^1[B]^1 or rate = k[A]^2 or rate = k[B]^2
where k has units M^-1s^-1
often suggests a physical collision between two reactant molecules, especially if the rate law is first-order with respect to each of the two reactants
why are there few processes with third-order rates?
it is far more rare for three particles to collide simultaneously with the correct orientation and sufficient energy to undergo a reaction
defn: mixed-order reactions
- non-integer orders (fractions) = broken-order
- reactions with rate orders that vary over the course of the reaction
be responsible for recognizing how the rate order changes as the reactant concentration changes
