Kinetics

Chemical Nature: Bond types and inherent molecular reactivity.
Contact (Surface Area):
- Homogeneous (same phase) = rapid mixing.
- Heterogeneous (different phases) = restricted to interface; higher surface area increases rate.
Concentration: More particles mean more collisions, increasing the rate.
Temperature: Higher thermal energy increases collision frequency and force. Rule of thumb: Rate doubles per 10°C rise.
Catalysts: Accelerate reactions significantly (up to x10^6 times) by offering an alternative path without being consumed.

Rate Change over Time: Reaction rate is not constant. It is initially fast (high reactant concentration) and gets slower as reactants are depleted.
Rate Law vs. Time Dependence:
- Rate Law: Connects reaction speed to concentrations.
- Time Dependence: Required to find exact concentrations at a specific time or to calculate how long it takes for a reactant to drop below a minimum value.

Reaction Orders and Rate Changes

0th order: No change to the rate. (No matter what you do to the concentration, the speed stays exactly the same).
1st order: Matches the change perfectly. (If you double concentration, rate doubles; if you triple it, rate triples).
2nd order: Squares the change. (If you double concentration, rate increases [2]² = 4 times; if you triple it, rate increases [3]^2 = 9 times).

Basis: Rate is proportional to effective collisions per second.
Requirements for a Reaction:
1. Molecules must collide.
2. Must have a minimum kinetic energy (KE).
3. Must be correctly oriented.
Concentration effect: Higher concentration = higher rate.
Temperature effect: Higher temp $\rightarrow$ higher molecular speed = higher fraction of collisions with enough force to react.

Activation Energy (Ea): Minimum KE required for a reaction to occur upon collision.
Transition State Theory: Visualizes collisions and the relationship between Ea and total potential energy (PE) via a PE diagram.
Exothermic (ΔHrxn<0): System PE decreases, KE/temperature increases (releases heat). Ea can be high or low regardless of ΔH.
Endothermic (ΔHrxn>0): System PE increases, KE/temperature decreases. Must add energy; Ea≥ΔHrxn.

Reaction Mechanism: The entire sequence of individual steps.
Multi-step Mechanism: Contains two or more elementary steps that combine to yield the net reaction.
Unimolecular
- What it is: Only one reactant molecule is involved.
- What happens: The molecule breaks apart or rearranges all by itself. (doesnt need another compound)
- Rate Law: Always 1st order because the speed depends entirely on how many of those single molecules are sitting around waiting to split.
  $\text{Rate}=k[A]$
Biomolecular
- What it is: Two reactant molecules are involved.
- What happens: The two molecules must physically crash into each other to react.
- Rate Law: Always 2nd order overall because the speed depends on the odds of those two things colliding.
  $\text{Rate} = k[A][B] \quad \text{or} \quad \text{Rate} = k[A]^2$
Elementary Process: A single step that occurs exactly as written.
- Rate Law Rule: For elementary steps only, the exponents in the rate law equal the stoichiometric coefficients of the reactants.
- Molecularity: Defined by the number of molecules participating in that specific elementary step.

A catalyst speeds up a reaction but never gets used up. It exits the reaction completely untouched.

NOTES FOR SOLVING:

For stoichiometric Ratio, the entire reaction shares an equal rate, the only thing we look at is the ratio for how something is consumed or produce
- Ex: C3H8 + 5O2 → 3CO2 + 4H2O

Just remember that reactant is Negative and coefficients must be reciprocal'd

This means ex: for every C3H8 consumed, 5 O2’s are consumed/ 3 CO2’s are produced

For integrated rate law, its function is to show how fast something is reacting. On the other hand integrate rate laws function is to show how much of something you have left after a specific amount of time