chemical kinetics (or reaction kinetics)
collision theory
reaction mechanism
activated complex (קומפלקס משופעל, תצמיד משופעל)
- transition state (מצב מעבר)
activation energy ( $E_{a}$ ) (אנרגיית שפעול)
effective collision (התנגשות פוריה)
catalysis (זֵירוז, קִטלוז, קטליזה)
- catalyst (זָרָז, קָטָלִיזָטוֹר)

Reaction rate

reaction rate (or rate of reaction) (קצב תגובה, מהירות תגובה)
- For $a A + b B p P + q Q$
  - $a, b, p, q$ are the stoichiometric coefficients
  - $v = - \frac{1}{a} \frac{d [ A ]}{d t} = - \frac{1}{b} \frac{d [ B ]}{d t} = \frac{1}{p} \frac{d [ P ]}{d t} = \frac{1}{q} \frac{d [ Q ]}{d t}$
- symbols: $v$ , $ν$ , $r$ , $R$ , $Rate$
- units: $M/s$ (molarity per second)
rate equation (or rate law or empirical differential rate equation)
reaction rate constant (or rate coefficient) $k$
- Arrhenius equation
an increase in temperature, surface area, the presence of a catalyst, and an increase in pressure (for gaseous reactants) generally increases the reaction rate

$a A products$

$v = k [A]^{n}$ , where:
- $n$ is the reaction order (determined experimentally. It is not necessarily $n = a$ unless the reaction is elementary)

	zero-order ( $n = 0$ )	first-order ( $n = 1$ )	second-order ( $n = 2$ )
differential rate law	$Rate = k$	$Rate = k [A]$	$Rate = k [A]^{2}$
integrated rate law	$[A]_{t} = - k t + [A]_{0}$	$ln [A]_{t} = - k t + ln [A]_{0}$	$\frac{1}{[ A ] _{t}} = k t + \frac{1}{[ A ] _{0}}$
alternative integrated form	$[A]_{0} - [A]_{t} = k t$	$[A]_{t} = [A]_{0} e^{- k t}$	$\frac{[ A ] _{0} - [ A ] _{t}}{[ A ] _{0} [ A ] _{t}} = k t$
unit of $k$	$M \cdot s^{- 1}$	$s^{- 1}$	$M^{- 1} \cdot s^{- 1}$
half-life ( $t_{1/2}$ )	$t_{1/2} = \frac{[ A ] _{0}}{2 k}$	$t_{1/2} = \frac{l n ( 2 )}{k} \approx \frac{0.693}{k}$	$t_{1/2} = \frac{1}{k [ A ] _{0}}$

$v = k [A]^{m} [B]^{n}$ where:
- $n$ and $m$ are the partial orders, and $m + n$ is the overall order
  - determined experimentally. It is not necessarily $n = b, m = a$ unless the reaction is elementary