kinetics

Rate:
The term kinetics implies motion. Kinetic energy is that of an object in motion. In chemistry, kinetics is the study of a chemical reaction "in motion", or in progress.

The kinetics of a chemical reaction revolve around its rate of reaction:


 * rate = -d[A]/dt**

where d[A] is the change in concentration, and dt is the change in time. The sign is negative due to the fact that the concentration will drop over time.

As an analogy, think of rate as speed: the rate of travel is equal to the change in distance over change in time.

Look at what happens in a hypothetical reaction A --> B:



The initial conditions (such as that in an ICE problem) are:
 * [A] is at its maximum value (known as [A] 0 )
 * [B] = 0

As time progresses from t = 0 to 80 seconds, changes will occur (C in ICE):
 * [A] is going to decrease (its rate of consumption will eventually slow)
 * [B] is going to increase (its rate of production will eventually slow)

Eventually, the reaction reaches equilibrium (E in ICE) at around 80 seconds:
 * the [A] and [B] will stabilize and reach a steady state (notice their lines are parallel in the plot above)
 * the rate of consumption of A is equal to the rate of consumption of B.

In studying equilibrium, the focus was on what happens after equilibrium sets in (i.e. after 80 sec above). In kinetics, the study focuses on the changes that occur from time zero to before equilibrium occurs. More specifically, the instantaneous changes that occur as soon as the reaction begins are of most importance.

Rate Laws
Another difference between kinetics and equilibrium is the mathematical expression used.

Kinetics only uses the reactants (equilibrium uses both) of a chemical equation. It also ignores stoichiometric coefficients. The mathematical expression is called the **rate law**.

For example, a reaction A --> B would have a rate law of:

rate = k[A] x


 * recall that rate is the change in concentration versus change in time (M/sec, for example);
 * the letter k is the rate constant;
 * x is called an order of reaction.

As an example: N 2 (g) + 3 H 2 (g) --> 2 NH 3 (g)

The rate law would be:

rate = k[N 2 ] x [H 2 ] y

where x and y are the reaction orders of nitrogen and hydrogen, respectively.

A **reaction order** describes the effect a change in concentration of that reactant has on the rate of the reaction.

The Collision Model
The success of a reaction between molecules is dependent on three things:
 * the molecules involved in the reaction;
 * the energy required to cause a reaction (called //activation energy//);
 * the orientation of the molecules during collision.

The following video explains some of these requirements.

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An example of molecular orientation is shown below:

the molecule with the yellow atom approaches the molecule with the purple atom, and they must be properly aligned. When they are, and have enough energy to react, a new (green) bond is formed, and one is broken (red). This process is usually sketched using an energy diagram.

There are a couple of important things to look at in an energy diagram:
 * 1) **Is it endo- or exothermic?** If the products have less energy than the reactants, it is exothermic, as shown in the diagram.
 * 2) **What is the activation energy?** This is the peak of the graph's curve, where the molecules are in the process of exchanging atoms.

Rate Changers A reaction rate can be influenced by a number of factors:
 * the composition of reactants (what they are made of, and what state of matter they are in)
 * the concentration of reactants
 * the temperature of the reaction mixture
 * the surface area of a solid
 * the agitation (mixing) of reactants
 * the presence of catalysts

A catalyst works by decreasing the energy barrier (activation energy) required by the reactants. Some catalysts hold a reactant in the proper alignment; others actually weaken some bonds, making them easier to break.

As shown in the energy diagram below, a catalyst will //lower the activation energy//:

Thus the rate of reaction will increase.

Here is a molecular model of a catalyst, holding a reactant (blue) in place at a better orientation for reaction: And here is a surface catalyst model: