Glossary
Activation energy
The minimum amount of energy required for a chemical reaction to proceed, representing the energy barrier that reactants must overcome to transform into products.
Example:
Enzymes in biological systems function by significantly lowering the activation energy for specific biochemical reactions, allowing them to occur rapidly at body temperature.
Collision model
A theoretical model that explains how chemical reactions occur, stating that reactant particles must collide with sufficient energy and the correct orientation.
Example:
The collision model helps us understand why increasing the concentration of reactants often increases the reaction rate, as it leads to more frequent molecular encounters.
Correct orientation
The specific spatial arrangement of colliding reactant molecules that allows their reactive parts to align properly, enabling the formation of new chemical bonds.
Example:
Just like a key needs the correct orientation to fit into a lock, reactant molecules must align precisely for an effective collision to occur.
Effective Collisions
Collisions between reactant particles that successfully lead to the formation of products because they meet both the energy and orientation requirements.
Example:
Only a small percentage of the countless molecular bumps in a reaction vessel are effective collisions that actually result in chemical change.
Enough energy
The minimum kinetic energy that colliding reactant particles must possess to overcome the activation energy barrier and successfully form products.
Example:
For a spark to ignite a flammable gas, it must provide enough energy to initiate the combustion reaction.
Ineffective Collisions
Collisions between reactant particles that do not result in product formation because they lack sufficient energy or the correct spatial orientation.
Example:
If two molecules collide too gently or at an awkward angle, it results in an ineffective collision, and they simply bounce off each other without reacting.
Kinetic Energy (Conservation Law)
A principle stating that the total kinetic energy of colliding particles remains constant before and after a collision, assuming no energy is converted to other forms.
Example:
In an ideal elastic collision between two gas molecules, their combined kinetic energy before the collision is equal to their combined kinetic energy after the collision.
Maxwell-Boltzmann distributions
Graphical representations that illustrate the range of kinetic energies or molecular speeds present in a sample of gas particles at a given temperature.
Example:
Analyzing Maxwell-Boltzmann distributions shows that at higher temperatures, a greater fraction of molecules possess the necessary activation energy for a reaction.
Momentum (Conservation Law)
A principle stating that the total momentum of a system of colliding particles remains constant, provided no external forces act on the system.
Example:
When a cue ball strikes another billiard ball, the total momentum of the system (both balls) is conserved, even as their individual velocities change.
Rate of chemical reactions
The measure of how quickly reactants are consumed and products are formed in a chemical reaction over a specific period.
Example:
The rate of chemical reactions for the decomposition of hydrogen peroxide can be sped up by adding a catalyst like potassium iodide.