Reaction rate is the change in concentration of reactants or products per unit time, typically expressed as [product]/t.

Pressure significantly affects gas-phase reactions but has minimal impact on liquids due to their incompressibility.

The order is the exponent in the rate law, so for Rate = k[A]^2, the order is 2.

A catalyst provides an alternative reaction pathway with lower activation energy, increasing the reaction rate.

For a first-order reaction, t_1/2 = ln(2)/k, which is constant and independent of initial concentration.

For a second-order reaction, Rate = k[A]^2, so units of k are (mol L^-1 s^-1)/(mol L^-1)^2 = L mol^-1 s^-1.

The Arrhenius equation, k = A e^(-Ea/RT), shows the dependence of the rate constant on temperature.

Since the rate law is first-order with respect to [A], doubling [A] doubles the rate (Rate ∝ [A]).

Activation energy is the energy barrier that reactants must overcome to form products.

A catalyst increases the reaction rate by lowering the activation energy, allowing more molecules to react.

For a zero-order reaction, Rate = k, which is constant regardless of reactant concentration.

For a second-order reaction, t_1/2 = 1/(k[A]_0), which depends on the initial concentration.

For a first-order reaction, the half-life is constant (t_1/2 = ln(2)/k), independent of concentration.

Higher temperatures increase the rate constant by providing more molecules with energy to overcome the activation barrier.

If rate doubles (2) when [A] quadruples (4), then Rate ∝ [A]^n, so 2 = 4^n, and n = log(2)/log(4) = 0.5.

Higher temperature increases molecular collisions and the fraction of molecules with energy above the activation energy, per the Arrhenius equation.

For heterogeneous reactions, surface area, concentration, and temperature affect the rate, but volume is less relevant unless it changes pressure for gases.

The rate law is determined experimentally by measuring how rate changes with concentration, not from stoichiometry or equilibrium.

For a first-order reaction, ln[A] = -kt + ln[A]_0, so plotting ln[A] vs. time yields a straight line.

A catalyst increases the reaction rate by lowering the activation energy without affecting equilibrium or yield.