Gen Chem II Exam 2

Chemical kinetics

Study of the rates of chemical reactions.

Reaction rate

Change in concentration of a reactant or product per unit time.

Rate sign for a reactant

Use a negative sign because reactant concentration decreases with time.

Rate sign for a product

Positive because product concentration increases with time.

Stoichiometric rate relationship

Divide each species' concentration change by its stoichiometric coefficient so all expressions give the same reaction rate.

Average reaction rate

Concentration change divided by the corresponding time interval.

Instantaneous reaction rate

Slope of the tangent to a concentration-versus-time curve at a particular time.

Why reaction rate usually decreases as a reaction proceeds

Reactant concentrations decrease, so effective collisions become less frequent.

Factors that affect reaction rate

Physical state/surface area, reactant concentration, temperature, and catalysts.

Surface area and reaction rate

Greater surface area of a solid generally increases reaction rate.

Concentration and reaction rate

Higher reactant concentration generally increases collision frequency and reaction rate.

Temperature and reaction rate

Higher temperature generally increases the reaction rate.

Collision theory

Reaction rate depends on the number of effective collisions per unit time.

Effective collision

A collision with proper orientation and sufficient energy to react.

Two requirements for an effective collision

Proper orientation and sufficient collision energy.

Activation energy (Ea)

Minimum energy barrier that must be overcome for reaction to occur.

If collision energy < Ea

The collision does not produce reaction.

If collision energy > Ea

The collision can produce reaction if orientation is also correct.

Transition state / activated complex

Highest-energy, unstable arrangement along the reaction pathway.

Reaction energy profile

Plot of energy versus reaction progress.

Forward activation energy

Energy difference between reactants and the transition state.

Reverse activation energy

Energy difference between products and the transition state.

Exothermic energy profile

Products are lower in energy than reactants; ΔH < 0.

Endothermic energy profile

Products are higher in energy than reactants; ΔH > 0.

Effect of larger Ea

Generally a slower reaction because fewer molecules can overcome the barrier.

Rate law

Equation relating reaction rate to reactant concentrations.

General rate law

Rate = k[A]^m[B]^n, where m and n are experimentally determined reaction orders.

Can reaction orders usually be taken from the balanced equation?

No. They must be determined experimentally, except for an elementary reaction.

Method of initial rates

Compare experiments in which one reactant concentration changes while the others remain constant.

Overall reaction order

Sum of the individual exponents in the rate law.

Zero order in A

Changing [A] has no effect on rate.

First order in A

Rate is directly proportional to [A].

Second order in A

Rate is proportional to [A]^2.

Doubling a zero-order reactant

Rate does not change.

Doubling a first-order reactant

Rate doubles.

Doubling a second-order reactant

Rate quadruples.

Tripling a first-order reactant

Rate triples.

Tripling a second-order reactant

Rate increases by 9 times.

Rate constant k

Proportionality constant in the rate law; depends on temperature and the reaction, not reactant concentration.

Units of k for a zero-order reaction

M·time^-1.

Units of k for a first-order reaction

time^-1.

Units of k for a second-order reaction

M^-1·time^-1.

Units of k for a third-order reaction

M^-2·time^-1.

Zero-order integrated rate law

[A]t = -kt + [A]0.

Linear plot for zero order

[A] versus time.

Zero-order plot slope

-k.

First-order integrated rate law

ln[A]t = -kt + ln[A]0.

Linear plot for first order

ln[A] versus time.

First-order plot slope

-k.

Second-order integrated rate law

1/[A]t = kt + 1/[A]0.

Linear plot for second order

1/[A] versus time.

Second-order plot slope

+k.

First-order half-life equation

t1/2 = 0.693/k.

First-order half-life and initial concentration

Independent of initial concentration.

Radioactive decay kinetics

Nuclear decay processes follow first-order kinetics.

Meaning of half-life

Time required for the amount/concentration of a reactant to fall to one-half its initial value.

Amount after 1 half-life

1/2 of the original amount.

Amount after 2 half-lives

1/4 of the original amount.

Amount after 3 half-lives

1/8 of the original amount.

Arrhenius equation

k = A e^(-Ea/RT).

Arrhenius equation: A

Frequency factor; reflects collision frequency/orientation probability.

Arrhenius equation: R

Gas constant; units must be consistent with Ea.

Temperature units in Arrhenius calculations

Kelvin.

Effect of increasing temperature on k

k increases.

Why higher temperature increases reaction rate

A larger fraction of molecules has enough kinetic energy to overcome Ea.

Reaction mechanism

Stepwise sequence of elementary reactions that produces the overall reaction.

Elementary reaction

A reaction that occurs in a single molecular event/step.

Molecularity

Number of reactant particles involved in an elementary step.

When can a rate law be written directly from coefficients?

For an elementary reaction.

Valid mechanism requirement #1

Elementary steps must add to the overall balanced reaction.

Valid mechanism requirement #2

Predicted rate law must agree with the experimentally determined rate law.

Rate-determining step

Slow step that controls the overall reaction rate.

Intermediate

Species produced in one step and consumed in a later step; absent from the overall reaction.

Catalyst in a mechanism

Species consumed in one step and regenerated in a later step; absent from the overall reaction.

Catalyst vs. intermediate memory trick

Catalyst: IN then OUT (consumed then regenerated). Intermediate: OUT then IN (produced then consumed).

Catalyst

Substance that increases reaction rate without being consumed overall.

How a catalyst speeds a reaction

Provides an alternative pathway with lower activation energy.

Does a catalyst change ΔH?

No.

Does a catalyst change ΔG?

No.

Does a catalyst change the equilibrium constant?

No.

Does a catalyst make a nonspontaneous process spontaneous?

No.

Homogeneous catalysis

Catalyst is in the same phase as the reactants.

Heterogeneous catalysis

Catalyst is in a different phase from the reactants.

Adsorption

Binding of particles to a surface.

Absorption

Uptake of particles into the interior of a substance.

Enzyme

Biological catalyst.

Active site

Region of an enzyme where the reaction occurs.

System

The matter under investigation.

Surroundings

Everything in the universe outside the system.

Universe in thermodynamics

System + surroundings.

Spontaneous process

Process that, once initiated under given conditions, can proceed without continuous outside intervention.

Nonspontaneous process

Process that does not occur on its own under the specified conditions.

Spontaneity vs. speed

Spontaneity is thermodynamic favorability; it does not tell how fast a reaction occurs.

Entropy (S)

Measure related to the number of possible microscopic arrangements (microstates) of a system.

Microstate

One possible arrangement of particle positions and energies.

Entropy and number of microstates

More available microstates → greater entropy.

Entropy and gas volume

Increasing the volume available to a gas generally increases entropy.

Motional energy types

Translational, rotational, and vibrational.

Molecular complexity and entropy

More complex molecules generally have more vibrational modes and higher molar entropy.

Standard molar entropy (S°)

Entropy of one mole of a substance under standard conditions; commonly J/(mol·K).