Geochemical Kinetics. Look at 3 levels of chemical change: Phenomenological or observational Measurement of reaction rates and interpretation of data in terms of rate laws based on mass action Mechanistic

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Geochemical Kinetics

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In it’s most general representation, we can discuss a reaction rate as a function of the extent of reaction:

Rate = dξ/Vdt

where ξ (small ‘chi’) is the extent of rxn, V is the volume of the system and t is time

Normalized to concentration and stoichiometry:

rate = dni/viVdt = d[Ci]/vidt

where n is # moles, v is stoichiometric coefficient, and C is molar concentration of species i

Rate Law

For any reaction: X Y + Z

We can write the general rate law:

Rate = change in concentration of X with time, t

Order of reaction

Rate Constant

Concentration of X

Reaction Order

ONLY for elementary reactions is reaction order tied to the reaction

The molecularity of an elementary reaction is determined by the number of reacting species: mostly uni- or bi-molecular rxns

Overall reactions need not have integral reaction orders – fractional components are common, even zero is possible

General Rate Laws

First step in evaluating rate data is to graphically interpret the order of rxn

Zeroth order: rate does not change with lower concentration

First, second orders:

Rate changes as a function of concentration

Zero Order

Rate independent of the reactant or product concentrations

Dissolution of quartz is an example:

SiO2(qtz) + 2 H2O H4SiO4(aq)

log k- (s-1) = 0.707 – 2598/T

First Order

Rate is dependent on concentration of a reactant or product

Pyrite oxidation, sulfate reduction are examples

First Order

Find order from log[A]t vs t plot

Slope=-0.434k

k = -(1/0.434)(slope) = -2.3(slope)

k is in units of: time-1

1st-order Half-life

Time required for one-half of the initial reactant to react

Second Order

Rate is dependent on two reactants or products (bimolecular for elementary rxn):

Fe2+ oxidation is an example:

Fe2+ + ¼ O2 + H+ Fe3+ + ½ H2O

General Rate Laws

2nd Order

For a bimolecular reaction: A+B products

[A]0 and [B]0 are constant, so a plot of log [A]/[B] vs t yields a straight line where slope = k2 (when A=B) or = k2([A]0-[B]0)/2.3 (when A≠B)

Pseudo- 1nd Order

For a bimolecular reaction: A+B products

If [A]0 or [B]0 are held constant, the equation above reduces to:

SO – as A changes B does not, reducing to a constant in the reaction: plots as a first-order reaction

2nd order Half-life

Half-lives tougher to quantify if A≠B for 2nd order reaction kinetics – but if A=B:

If one reactant (B) is kept constant (pseudo-1st order rxns):

3rd order Kinetics

Ternary molecular reactions are more rare, but catalytic reactions do need a 3rd component…

Zero order reaction

NOT possible for elementary reactions

Common for overall processes – independent of any quantity measured

[A]0-[A]=kt

Pathways

For an overall reaction, one or a few (for more complex overall reactions) elementary reactions can be rate limiting

Reaction of A to P rate determined by slowest reaction in between

If more than 1 reaction possible at any intermediate point, the faster of those 2 determines the pathway

Initial Rate, first order rxn example

For the example below, let’s determine the order of reaction A + B C

Next, let’s solve the appropriate rate law for k

Rate Limiting Reactions

For an overall reaction, one or a few (for more complex overall reactions) elementary reactions will be rate limiting

Reaction of A to P rate determined by slowest reaction in between

If more than 1 reaction possible at any intermediate point, the faster of those 2 determines the pathway

Activation Energy, EA

Energy required for two atoms or molecules to react

Transition State Theory

The activation energy corresponds to the energy of a complex intermediate between product and reactant, an activated complex

A + B ↔ C± ↔ AB

It can be derived that

EA = RT + DHC±

Collision Theory

collision theory is based on kinetic theory and supposes that particles must collide with both the correct orientation and with sufficient kinetic energy if the reactants are to be converted into products.

The minimum kinetic energy required in a collision by reactant molecules to form product is called the activation energy, Ea.

The proportion of reactant molecules that collide with a kinetic energy that is at least equal to the activation energy increases rapidly as the temperature increases.

T dependence on k

Svante Arrhenius, in 1889, defined the relationship between the rate constant, k, the activation energy, EA, and temperature in Kelvins:

or:

Where A is a constant called the frequency factor, and e–EA/RT is the Boltzmann factor, fraction of atoms that aquire the energy to clear the activation energy

Arrhenius Equation

y = mx + b

Plot values of k at different temperatures: log k vs 1/T slope is EA/2.303R to get activation energy, EA

Activation Energy

EA can be used as a general indicator of a reaction mechanism or process (rate-limiting)