CHEMICAL KINETICS
· area of Chemistry concerned with the speed or rates at which a chemical reaction occurs.
· “kinetic” suggests movement or change.
· Refers to the rate of a reaction which is the change in concentration of a reaction on a product with time
Reaction A-->product OR simply A --> B
§ Reaction Rates
*ratereactant=D[A]
D t
-the lower the rate the lower the concentration
the higher the rate the higher the concentration
§ FACTORS AFFECTING THE RATE OF CHEMICAL REACTION
1. 1. Temperature- Reaction rate increases with and increase in temperature. This happens because the kinetic energy of the reactants increase with a rise in temperature.
2. 2.Concentration of Reactants- An increase in concentration means an increase in the molecules or particles per unit volume. And thus a decrease in space between the reacting particles with less distance to travel inside the vessel, the more the frequent the collision, the faster the reaction.
3. 3.Presence of catalyst- Catalyst-a substance that speeds up a chemical reaction without itself undergoing a chemical change.
-Inhibitors- substance that slows down chemical reaction.
-Enzymes-catalyst in organisms.
4. 4. Nature of Reaction
Ionic compound-reacts easily
Covalent compound-reacts slowly
Polar-polar- mixes or dissolves easily/ same with Non-polar
Homogenous reactions- reacts faster than the heterogenous
5. 5.Surface area -the greater the surface area the faster the rate of reaction.
-the smaller the size of particles the larger the surface area.
§ THE COLLISION THEORY OF CHEMICAL KINETICS
-Collision theory views the reaction rate as the result of the particles collidng with a certain frequency and minimum energy.
-states that molecules must collide in order to react.
*not all collisions are effective.
Effective Collision- leads to the formation of products.
Ineffective Collision-does not lead to the formation of products.
Two Factors Affecting Strong Collision
· Sufficient energy
· Proper orientation of colliding particles
Activation Energy-minimum energy required to activate the molecules
Activated Complex (Transition State)-formed when molecules collide.
-a temporary species formed by the reactant molecules as a result of the collision before they form the product.
§ THE RATE LAW
The rate law expresses the relationship of the rate of a reaction to the rate constant and the concentrations of the reactants.
aA + bB à cC + dD **small letter-corresponds to coefficient
big letter- substance involved
where [A] and [B] express the concentration of the species A and B, respectively (usually in moles per liter (molarity)); x and y are not the respective stoichiometric coefficients of the balanced equation; they must be determined experimentally. k is the rate coefficient or rate constant of the reaction. The value of this coefficient k depends on conditions such as temperature, ionic strength, surface area of the adsorbent or light irradiation.
coefficients of the balanced equation; they must be determined experimentally. k is the rate coefficient or rate constant of the reaction. The value of this coefficient k depends on conditions such as temperature, ionic strength, surface area of the adsorbent or light irradiation.
Differential Rate Laws
In many reactions, the rate of reaction changes as the reaction progresses. Initially the rate of reaction is relatively large, while at very long times the rate of reaction decreases to zero (at which point the reaction is complete). In order to characterize the kinetic behavior of a reaction, it is desirable to determine how the rate of reaction varies as the reaction progresses.A rate law is a mathematical equation that describes the progress of the reaction. In general, rate laws must be determined experimentally. Unless a reaction is an elementary reaction, it is not possible to predict the rate law from the overall chemical equation. There are two forms of a rate law for chemical kinetics: the differential rate law and the integrated rate law.
The differential rate law relates the rate of reaction to the concentrations of the various species in the system.
Differential rate laws can take on many different forms, especially for complicated chemical reactions. However, most chemical reactions obey one of three differential rate laws. Each rate law contains a constant, k, called the rate constant. The units for the rate constant depend upon the rate law, because the rate always has units of mole L-1 sec-1 and the concentration always has units of mole L-1.
Zero-Order Reaction
For a zero-order reaction, the rate of reaction is a constant. When the limiting reactant is completely consumed, the reaction abrupts stops.First-Order Reaction
Differential Rate Law: r = k
The rate constant, k, has units of mole L-1 sec-1.
For a first-order reaction, the rate of reaction is directly proportional to the concentration of one of the reactants.Second-Order Reaction
Differential Rate Law: r = k [A]
The rate constant, k, has units of sec-1.
For a second-order reaction, the rate of reaction is directly proportional to the square of the concentration of one of the reactants.
Differential Rate Law: r = k [A]2
The rate constant, k, has units of L mole-1 sec-1.
REACTION RATES AND STOICHIOMETRY
Consider the reaction. 4NO2 (g) + O2 (g) --> 2N2O5(g)
Suppose that at a particular moment during the reaction, molecular oxygen is reacting at a rate of 0.024 m/s.
a) At waht rate is N2O5 being formed?
b) At what rate is NO2 reacting?
a. a. Rate oxygen= 0.024 m/s
1 [N2O5] = d[O2] = 0.024 m/s
2 dt dt
b.
Rate Data for the Reaction between F2 and ClO2
[F2] (M) | [ClO2] (M) | Initial Rate (m/s) |
1 0.10 | 0.010 | 1.2 x 10 -3 |
2 0.10 | 0.040 | 4.8 x 10 -3 |
3 0.20 | 0.010 | 2.4 x 10 -3 |
Rate= K[F2]x [ClO2] y
X= log rate ratio
Log concentration ratio
Log rate ratio= log2.4 x10 –3 m/s = 2
log1.2 x 10 -3 m/s
Log concentration ratio= Log0.20 M =2
Log 0.10 M
**x=2/2=1, 1st order [F2]
Y= log rate ratio
Log concentration ratio
=Log 4.8x10-3m/s =4
Log 1.2x10 -3m/s
= Log0.040 M =4
Log 0.010 M
**y=4/4= 1
§ REACTION MECHANISM
-sequence of elementary steps that leads to product formation.
-complete description of how the reactants are converted to product.
A mechanism describes in detail exactly what takes place at each stage of a chemical transformation. It also describes each transition state, which bonds are broken (and in what order), which bonds are formed (and in what order) and what the relative rates of the steps are. A complete mechanism must also account for all reactants used, the function of a catalyst, stereochemistry, all products formed and the amount of each.
A reaction mechanism must also account for the order in which molecules react. Often what appears to be a single step conversion is in fact a multistep reaction.
Examples
Consider the following reaction:
CO + NO2 → CO2 + NO
In this case, it has been experimentally determined that this reaction takes place according to the rate law R = k[NO2]2. Therefore, a possible mechanism by which this reaction takes place is:
2 NO2 → NO3 + NO (slow)
NO3 + CO → NO2 + CO2 (fast)
Each step is called an elementary step, and each has its own rate law and molecularity. All of the elementary steps must add up to the original reaction, by means of organic reactions (e.g.: rearrangement reactions).
When determining the overall rate law for a reaction, the slowest step is the step that determines the reaction rate. Because the first step (in the above reaction) is the slowest step, it is the rate-determining step. Because it involves the collision of two NO2 molecules, it is a bimolecular reaction with a rate law of R = k[NO2]2. If we were to cancel out all the molecules that appear on both sides of the reaction, we would be left with the original reaction.
§ ELEMENTARY PROCESS OR REACTION
-series of single-step changes
-also known as elementary steps or series of simple reactions that represent the progress of the overall reaction.
§ MOLECULARITY OF A REACTION
-is the number of molecules reacting in an elementary steps.
A reaction involving one molecular entity is called unimolecular.
A reaction involving two molecular entities is called bimolecular.
A reaction involving three molecular entities is called termolecular
A good example is the reaction of CO and NO2, CO(g) + NO2(g) -> CO2(g) + NO(g). At low temperature, this reaction occurs in two steps
NO2(g)+ NO2(g) -> NO3(g) + NO(g) +
CO(g) + NO3(g) -> CO2(g) + NO2(g) =
CO(g) + NO2(g) -> CO2(g) + NO(g)
The first two reactions are the two elementary steps in this reaction.
Each elementary step proceeds at a different rate. In the above equation, the first reaction is slow and the second is very fast. This means that the rate of the overall reaction is dominated by the rate of the first reaction: this is the rate determining step. The rate expression for the overall reaction is
k = [NO2]2
since the first elementary step is the rate determining step.
The rate expression for a complex reaction can be determined by looking at the overall steps: simply write down the elementary steps, figure out which is the slowest, then write the reaction rate expression in terms of that step. Occasionally, however, a reactive intermediate can be important in the reaction: these are difficult to deal with since their concentration is very low and there is usually no easy way to measure it, thus making the reaction rate expression not very useful. We can remove the intermediate term by noting that the fast reactions that form reactive intermediates typically quickly establish an equilibrium: we can then express the concentration of the reactive intermediate in terms of things in the fast reaction.
This sounds more complex than it is: conside the reaction of NO and Cl2 to form NOCl. This reaction occurs in two elementary steps: the first is fast, the second is slow.
NO(g) + Cl2(g) < = > NOCl2(g) (fast)
NOCl2(g) + NO -> 2NOCl (slow)
2NO(g) + Cl2 -> 2NOCl
We can write the rate expression as
rate = k2[NOCl2][NO]
however, NOCl2 is a reactive intermediate with a concentration we probably can't measure. We can remove it by noting that the first reaction quickly establishes an equilibrium: the rate of the forward reaction (rate = k[NO][Cl2]) is equal to the rate of the reverse reaction (rate = k-1[NOCl2]). Therefore
k[NO][Cl2] = k-1[NOCl2]
[NOCl2] = k1[NO][Cl2]/k-1
Now we can just substitute this into our original rate expression to find
rate = k2[NO][NOCl2]
rate = k2[NO]k1[NO][Cl2]/k-1
rate = k2k1[NO]2[Cl2]/k-1
Example: The reaction 2NO(g) + O2(g) ->NO2(g) possibly has the following mechanism:
NO + O2 < = > NO3 (fast)
NO3 + NO - > 2NO2 (slow)
What is the rate expression for this reaction?
Solution: the basic rate expression is dependant on the rate of the second reaction, so the rate for the overall reaction is
rate = k2[NO3][NO]
However, NO3 is a reactive intermediate and we must remove it from the reaction rate expression. To do this, note that the first reaction rapidly establishes an equilibrium and the forward and backwards rates are identical. Thus
k1[NO][O2] = k-1[NO3]
[NO3] = k1[NO][O2]/k-1
We can now substitute this into our original rate expression
rate = k2[NO3][NO]
rate = k2[NO] k1[NO][O2]/k-1
rate = k2k1[NO]2[O2]/k-1
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