Introduction
The collision theory is essentially a detailed section of the kinetic-molecular theory; it is strongly interconnected to chemical kinetics. The theory is basically used to explain how different variables change the rate of reaction. This theory is based on the idea that reactant particles must collide for a reaction to occur, but only a small percentage of the total collisions have the energy and "orientation" to connect effectively and cause the reactants to form into the products. The field of study involving molecular collisions developed vigorously in the 1980’s because of modern electronic and vacuum instrumentation. Although the collision theory was proposed by Max Trautz and William Lewis in the early 1900's.
Discovering the Collision Theory:
Collision frequency-This is the count of how many times a particular molecule collides with others per unit of time. To create a reaction, species need to collide. All collisions do not create a chemical reaction. The atoms for instance need to collide orientated in a specific way, and they need to collide with a precise energy in order for bonds to break.
Effects of Orientation
In any collision with unsymmetrical atoms, the oritentation of the atoms during collision are crucial to creating a reaction. The reaction rate is effected by the orientation because without the proper orientation during collision, the reaction will notoccur at all.
Ex 1. A collision between two molecules: ethene, CH2=CH2, and hydrogen chloride, HCl. These react to give chloroethane. Because of the collision, the double bond in the middle of the two carbons is changed into a single bond. A hydrogen atom is attached to one of the carbons and a chlorine atom is attached to the other. The absolute only way that the reaction can happen is if the hydrogen side of the H-Cl bond meets the carbon-carbon (double bond).
Figure 1: Click to view larger.
Ex 2. When two Hydrogen atoms combine to create a hydrogen molecule--no bonds are destroyed. (An H-H bond is created) Hydrogen are all spherical and symmetrical in every way; this means that any way they decide to approach eachother is equal. The reaction would occur as quick as the collisions occur.
Figure 2: Click to view larger.
Effects of energy
Regardless of if the atoms are orientated in the correct way, a reaction will still not be able to occur, unless the particles collide with the activation energy of the reaction.
Activation energy- Is the minimum amount of energy above the average kinetic energy needed for a reaction to occur.
If the atoms collide with less energy than the, called for, activation energy--nothing happens at all. The atoms simply bounce off and away from eachother. Only the collisions that have energies equal to or greater than the activation energy create a reaction. Ultimtely, chemical reactions involve the breaking of some bonds, which takes energy, and making new bonds, which releases energy. Activation energy is the key to breaking the initial bonds. When collisions are too gengtle, the adequate amount of energy is not brought to the bonds-- a reactions fails to occur. The Activation energy is crucial in the reaction rate because depending on how much kinetic energy is brought to the collision-- the reaction will vary in speed and frequency.
Maxwell-Boltzmann Dsitribution
It is very useful to know the types of proportion of atoms that have high enough energies to cause a chemical reaction, when colliding. Gases can be plotted on a graph called the Maxwell-Boltzmann Distribution. It shows different atoms or particles, and their energies.
Figure 3: Click to view larger.
- Most of the particles have a moderate amount of energy.
- Only a select few have very high energies.
- Few also have very low energies.
Figure 4: Click to view larger.
-At higher temperatures, more molecules have higher kinetic energies.
References
- Collision Theory, Goldberg, Watson, 1964
- Atomic and Molecular Collison Theory, Gianturco, 1980
- General Chemistry, Principles & modern Applications,Petrucci,2007,2002, 1997
Problems
Example #1: In a classroom there are a group of blindfolded students and you have them walk around the room at a slow pace. A pair of students will bump into one another on occassion. If the students start to run around the room at a higher pace (higher pace = more energy for collision), then a collision is most likely going to be successful. If a portion of students move around fast, while others move about at a more sluggish pace, successful reactions will still occur, but not as often as if all students are running. Do you see the concept of energy and speed, and how they directly effect a successful or unsuccessful collision?
Example #2: The students are walking around at a pretty high speed but not running. If an arm to arm collision happens, then this is considered an unsuccessful reaction. On the other hand if one student steps on another's shoe, a successful collision occurs. You will see that of the collisions that happen, only a few will be stepping on another's toes. This shows you that collisions must occur with the correct orientation. Many collisions between students will occur, but only a select few will be successful!
Example #3: Which beaker is going to have more collisions, thus more chance of reactivity?
Beaker A: 10 ml beaker filled with 10 ml of a mixture.
Beaker B: 20 ml beaker filled with 10 ml of the same mixture.
Solution: Beaker A is the correct answer because the concentration is increased with less space, thus giving the atoms a higher chance of having enough energy and hitting at the correct orientation.
Example #4: When you add vinegar to baking soda does the collision theory take place?
Solution: Yes, the fizzing and bubbling of the reaction requires the collision theory to occur in every facet. This is a real-life example of the collision theory.
Example #5:When using a pipet to distribute HCL into an ice bath of Mg and boiling bath of Mg-- which reaction will occur faster?
Solution: The boiling Mg will push the reaction to occur more quickly because ,remember, higher heat tends to result in higher kinetic energy.
It is not so clear to me from your text how collision theory relate all those things you to the reaction rate tough.
I kind of agree with professor Madsen though. Maybe you can enter in a paragraph about how the orientation of the collisions of particles affect the rate of a reaction. Other than that, I love the pictures you found to put in.
A couple suggestions, for the activation energy if you could try an example involving reactants and products, probably a graph to show how much energy is the minimum for the reaction to react.