Reaction Rate
Reaction rate is how quickly reactants are used up or products form; temperature, concentration, surface area, and catalysts all change it.
Definition
Reaction rate describes how fast a chemical reaction happens, ranging from a firecracker's split-second explosion to an iron nail's months-long rusting, and it is formally the change in concentration of a reactant or product per unit time, $\text{rate} = -\frac{1}{a}\frac{d[A]}{dt}$ for a reactant $A$ with coefficient $a$. According to collision theory, reactions occur when particles collide with enough energy and the right orientation, so rate is tied to concentrations through an experimentally determined rate law, $\text{rate} = k[A]^m[B]^n$, whose orders reflect the mechanism rather than the stoichiometry. The rate constant $k$ depends on temperature through the Arrhenius equation, $k = Ae^{-E_a/RT}$, so reactions with a higher activation energy are more sensitive to temperature change, and the slowest elementary step in a multistep mechanism is the rate-determining step.
Example
A whole antacid tablet fizzes in water for a minute or more, but the same tablet crushed to powder fizzes far faster because so much more surface is exposed, and warming the water speeds it up further. For a reaction with $E_a = 50\text{ kJ/mol}$, raising the temperature from $25^\circ\text{C}$ to $35^\circ\text{C}$ increases $k$ by a factor of about $1.9$, roughly the origin of the old rule that reaction rate doubles for every $10^\circ\text{C}$ rise.
Key Insight
Grain elevators and coal mines can explode because finely divided dust has enormous surface area, so the same material that burns slowly as a solid log or lump can react explosively as airborne powder. Thermodynamics and kinetics are entirely independent: diamond converting to graphite is spontaneous ($\Delta G < 0$) but immeasurably slow, and some highly exothermic reactions never start without a catalyst, so a favorable $\Delta G$ tells you nothing about how fast a reaction will go.