Ch. 5: Chemical Kinetics

Question 1

what does change in Gibbs free energy (delta G) denote?

Answer 1

determines whether or not a reaction will occur by itself without outside assistance (spontaneous or not)

Question 2

defn: reaction mechanism

Answer 2

proposed pathways/ the series of steps that proceed in order for a reaction to occur

sum of the steps gives the overall reaction

Question 3

defn: rate-determining step

Answer 3

the slowest step in any proposed mechanism

acts like a kinetic bottleneck, preventing the overall reaction from proceeding any faster than that slowest step

Question 4

defn + char + implications: collision theory of chemical kinetics

Answer 4

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

Question 5

defn + aka: activation energy (Ea)

Answer 5

part of the collision theory of chemical kinetics

the minimum energy of collision necessary for a reaction to take place

aka: energy barrier

Question 6

eqn: rate of a rxn (collision theory of chemical kinetics)

Answer 6

rate = Z x f

Z = total number of collisions occurring per second

f = the fraction of collisions that are effective

Question 7

eqn + defn: Arrhenius equation (collision theory of chemical kinetics) + defn: frequency factor

Answer 7

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)

Question 8

what relationships can you infer from the Arrhenius equation? (k = Ae^(-Ea/RT)

Answer 8

As A increases, k increases

reaction rate increases with temperature because exponent becomes smaller (becomes less negative)

Question 9

how can the frequency factor (A) be increased?

Answer 9

by increasing the number of molecules in a vessel

more molecules = more opportunities for collision

Question 10

expl: transition state theory

Answer 10

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

Question 11

defn: reaction coordinate (transition state theory)

Answer 11

traces the reaction from reactants to products, through the transition state

Question 12

defn + aka + char: transition state

Answer 12

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

Question 13

what is the difference between transition states and reaction intermediates?

Answer 13

transition states are theoretical constructs that exist at the point of max energy, rather than reaction intermediates which are distinct identities with finite timelines

Question 14

defn: free energy diagram

Answer 14

illustrates the relationship between the activation energy, the free energy of the reaction, and the free energy of the system

Question 15

defn: free energy of the reaction (deltaGrxn)

Answer 15

the difference between the free energy of the products and the free energy of the reactants

Question 16

defn: exergonic and endergonic reaction (in terms of free energy)

Answer 16

EXERGONIC REACTION = a negative free energy change // energy is given off

ENDERGONIC REACTION = a positive free energy change // energy is absorbed

Question 17

where is the transition state in the energy diagram?

Answer 17

at the peak of the energy diagram

Question 18

what are the four factors affecting reaction rate in chemical kinetics?

Answer 18

1. reaction concentrations
2. temperature
3. medium
4. catalysts

Question 19

how does reaction concentration affect reaction rate?

Answer 19

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

Question 20

how does temperature affect reaction rate?

Answer 20

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

Question 21

what temperature is an enzymatic reaction optimal at? what happens after this optimal point?

Answer 21

35 - 40 deg. C (body temperature)

after 40 deg, the curve falls sharply because denaturation has occurred

Question 22

are polar or nonpolar solvents generally preferred in terms of reaction rate?

Answer 22

polar –> bc the molecular dipole tends to polarize the bonds of the reactants, thereby lengthening and weakening them, allowing the rxn to occur faster

Question 23

defn + char: catalyst

Answer 23

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

Question 24

how can catalysts affect reaction rate?

Answer 24

1. increase the frequency of collisions between the reactants
2. change the relative orientation of the reactants (making higher % of the collisions effective)
3. donate electron density to the reactants
4. reduce intramolecular bonding within reactant molecules

Question 25

defn: homogeneous vs. heterogeneous catalysis

Answer 25

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

Question 26

what impact does a catalyst have on a reaction? what impact DOESN’T a catalyst have on a reaction?

Answer 26

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

Question 27

defn + eqn + units: rate for the general reaction aA + bB –> cC + dD

Answer 27

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

Question 28

what is the first thing to look for whenever a question asks to determine the rate law for a reaction?

Answer 28

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)

Question 29

eqn + unit: rate law for the general reaction aA + bB –> cC + dD

Answer 29

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

Question 30

4 common traps to pay attention to in chemical kinetics

Answer 30

1. 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)

2. 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

3. 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

4. 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

Question 31

how to use provided data in the experimental determination of rate law?

Answer 31

1. 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
2. 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
3. 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
4. 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
5. 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

Question 32

defn: zero-order reaction

Answer 32

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

Question 33

what are the two ways to change the rate of a zero-order reaction?

Answer 33

1. temperature
   - the rate constant itself is dependent on temperature
2. addition of a catalyst
   - lowers the activation energy, thereby increasing the value of k

Question 34

defn + implication: first-order reaction

Answer 34

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

Question 35

what is a classic example of a first-order reaction? + eqn

Answer 35

radioactive decay

[A]t=[A]0e^(-kt) = the concentration of radioactive substance A at any time t

Question 36

defn + implications: second-order reaction

Answer 36

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

Question 37

why are there few processes with third-order rates?

Answer 37

it is far more rare for three particles to collide simultaneously with the correct orientation and sufficient energy to undergo a reaction

Question 38

defn: mixed-order reactions

Answer 38

1. non-integer orders (fractions) = broken-order
2. 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