Glossary
Activated Complex (Transition State)
The highest energy, unstable intermediate structure formed at the peak of the potential energy diagram, where old bonds are breaking and new bonds are forming.
Example:
Imagine a brief, fleeting moment where a molecule is neither fully reactant nor fully product; that's the activated complex.
Activation Energy (Ea)
The minimum amount of energy required for reactants to transform into products, representing the energy difference between the reactants and the transition state.
Example:
To start a campfire, you need to provide enough heat to overcome the wood's activation energy and initiate combustion.
Arrhenius Equation
An equation that describes the temperature dependence of reaction rates, relating the rate constant to activation energy, temperature, and a frequency factor.
Example:
While you won't perform calculations with it on the AP exam, understanding the Arrhenius equation helps explain why reactions speed up when heated.
Bond Breaking
The process where chemical bonds between atoms are severed, which typically requires an input of energy.
Example:
During the combustion of methane, C-H and O=O bond breaking must occur before new bonds can form.
Bond Formation
The process where new chemical bonds are created between atoms, which typically releases energy.
Example:
The creation of water from hydrogen and oxygen involves H-O bond formation, which releases a significant amount of energy.
Elementary Reaction
A single-step reaction involving one or a few molecules, representing the basic building blocks of more complex reactions.
Example:
The decomposition of ozone, O₃ → O₂ + O, is an elementary reaction because it occurs in one step.
Endothermic Reactions
Chemical reactions that absorb energy from their surroundings, resulting in products with higher potential energy than reactants.
Example:
The dissolution of ammonium nitrate in water feels cold because it is an endothermic reaction absorbing heat from the surroundings.
Exothermic Reactions
Chemical reactions that release energy into their surroundings, resulting in products with lower potential energy than reactants.
Example:
The burning of natural gas is an exothermic reaction, releasing heat and light that can be used for cooking or heating.
First-order (reaction)
An elementary reaction whose rate depends linearly on the concentration of only one reactant.
Example:
If the rate of a reaction doubles when the concentration of reactant A doubles, it's likely a first-order reaction with respect to A.
Frequency Factor (A)
A term in the Arrhenius equation that accounts for the frequency of collisions and the probability that collisions have the correct orientation.
Example:
The frequency factor helps explain why some reactions are inherently faster due to more effective molecular encounters.
Gas Constant (R)
A fundamental physical constant used in many equations, including the ideal gas law and the Arrhenius equation, relating energy to temperature.
Example:
The gas constant (8.314 J/(mol·K)) is a universal value that connects energy and temperature in many chemical contexts.
Potential Energy Diagrams
Graphical representations that illustrate the change in potential energy of a system as a chemical reaction progresses from reactants to products.
Example:
A chemist might sketch a potential energy diagram to visualize the energy barrier and overall energy change for a new synthesis reaction.
Products
The new chemical substances formed as a result of a chemical reaction.
Example:
In the reaction 2H₂ + O₂ → 2H₂O, water is the product.
Rate Constant (k)
A proportionality constant in the rate law that relates the rate of a reaction to the concentrations of reactants at a given temperature.
Example:
A large rate constant indicates a fast reaction, while a small one suggests a slow reaction.
Reactants
The starting chemical substances that undergo a transformation during a chemical reaction.
Example:
In the reaction 2H₂ + O₂ → 2H₂O, hydrogen and oxygen are the reactants.
Second-order (reaction)
An elementary reaction whose rate depends on the concentration of two reactants, or the square of one reactant's concentration.
Example:
A reaction where two NO molecules collide to form N₂O₂ is a second-order elementary reaction, as its rate depends on [NO]².
Temperature (T)
A measure of the average kinetic energy of the particles in a substance, which directly influences reaction rates.
Example:
Increasing the temperature of reactants typically speeds up a reaction because molecules collide more frequently and with greater energy.