Rates of Reactions

rate

A measure of change over time.

particle theory

All matter is in constant motion with the speed represented by the average kinetic energy (temperature) of the system.

kinetic energy

Energy associated with the __motion__ of molecules.

collision theory

All particles in motion will randomly collide with one another. For a collision to result in a **reaction**, the collision must be **effective**. An effective collision (one that results in the **formation of products**) must satisfy **two** conditions.

effective collision

1. Reactant **orientations** must be favourable.
2. Collision must occur with sufficient **energy**.

Maxwell-Boltzmann Curve

### *A.K.A. Kinetic Energy Distribution:*

In most reactions, only a **small fraction** of the total collisions have sufficient energy for a reaction to occur. The collision energy depends on the **kinetic energy** of the colliding particles.

Plotting the number of particles versus the kinetic energy of the particles gives a curve like the one below:

**NOTE:** The curve represents the distribution of the kinetic energy of collisions __at a given temperature__. The activation energy is __**independent**__ of temperature (it does not change when temperature changes!).

potential energy diagram

Shows the relative **potential energies** of reactants, transition states and products as a reaction progresses.

potential energy

Energy __within__ molecules (i.e. the bonding energy).

activation energy

The __minimum__ amount of __kinetic energy__ that the reactants must collide with to make the reaction proceed (energy __barrier__ the reactants must overcome to form products).

* E*a*

activated complex/transition state

A __temporary__ arrangement of atoms that form as bonds are breaking and new bonds are forming (highly __unstable__).

endothermic

exothermic

surface area

When increasing this (**grinding, slicing, dissolving**) of __**solid**__ reactants, it exposes more molecules to each other than before, hence increasing the frequency/probability of effective collisions, increasing the rate (opposite occurs when this is *small*).

concentration

With an increase in the number of **reactant molecules present**, there is an increase in the frequency/probability of effective collisions occurring, hence increasing the rate (the reverse is true when this is *low*).

temperature

By increasing this, the molecules have more **kinetic energy** which increases the frequency/probability of effective collisions. With more energy, the **activation barrier** can be more easily overcome more often, increasing the rate (the *opposite* occurs when this is *low*).

volume/pressure

### ***** ONLY FOR A GAS SYSTEM!!! *****

By decreasing this factor of a container and and increasing the other, the molecules have less space to move and are in closer proximity to one another, hence increasing the frequency/probability of effective collisions, increasing the rate (*increasing* the former and *decreasing* the latter decreases the rate).

catalyst

Allows for an alternate pathway of **lower activation energy** so that more collisions can overcome the barrier more often, increasing the rate. They are **not consumed** in reactions and can be **reused**.

mol/Ls

Overall order: 0

(write units of the rate constant k)

/s

Overall order: 1

(write units of the rate constant k)

L/mols

Overall order: 2

(write units of the rate constant k)

L^2/mol^2s

Overall order: 3

(write units of the rate constant k)

rate law equation

Relates the **rate of a reaction** to the **concentrations of the reactants**, each concentration is expressed with an **order (exponent)**.

* **m** and **n** are the **reactant orders** determined from __experimental data__
* The **rate constant** is **k**, which makes the two sides of the equation equal to each other (units of k will change depending upon the __overall order__ of the reaction)

reaction mechanism

The overall sequence of **elementary steps** by which a chemical reaction occurs.

* A reaction that occurs in two or more elementary steps is called a multistep or complex reaction.

rate-determining step

When a reaction takes place in a series of steps, the overall rate of the chemical reaction is only as fast as the rate of its **slowest step** (the step with the **largest activation energy**).

* **The coefficient of each reactant in this is the order of each reactant in the rate law equation!**

reverse endothermic

reverse exothermic

reaction intermediates

Stable chemical compounds that do NOT appear in the overall reaction.

rate-determining step

The overall rate of the chemical reaction is only as fast as the rate of its slowest step (the step with the **largest activation energy**). The rate law equation is dependent on this step; **the coefficient of each reactant in the this step is the order of each reactant in the rate law equation!**